THE  4BSOEPTION  SPECTEA  OF  SOLUTIONS  AS  STUDIED 
BY  MEANS  OF  THE  EADIOMICEOMETEB^ 

THE  CONDUCTIVITIES,  DISSOCIATIONS,  AND  VISCOS- 
ITIES OF  SOLUTIONS  OF  ELECTEOLYTES 
IN  AQUEOUS,   NON-AQUEOUS,  AND 
MIXED  SOLVENTS 


BY 


HARRY  C.  JONES  AND  COLLABORATORS 


WASHINGTON,  D.  C. 

PUBLISHED  BY  THE  CARNEGIE  INSTITUTION  OF  WASHINGTON 
1915 


CARNEGIE  INSTITUTION  OF  WASHINGTON 
PUBLICATION  No.  210 


PRESS  OF  GIBSON   BROTHERS 
WASHINGTON,  D.  C. 


PREFACE. 

The  work  recorded  in  this  monograph,  while  apparently  dealing  with 
several  subjects,  is  in  fact  closely  connected,  in  that  it  all  bears  directly 
or  indirectly  on  the  present  solvate  theory  of  solution,  which  was  pro- 
posed in  the  Johns  Hopkins  laboratory  about  fifteen  years  ago. 

The  work  on  the  absorption  spectra  of  solutions  by  Dr.  Shaeffer  and 
Mr.  Paulus,  using  the  raoliomicrometer,  led  to  results  of  the  same  gen- 
eral character  as  those  obtained  earlier  by  Dr.  Guy  and  recorded  in 
publication  No.  190  of  the  Carnegie  Institution  of  Washington.  Solu- 
tions of  some  non-hydrated  salts  are  about  equally  transparent  with 
pure  water,  except  at  the  bottoms  of  the  absorption  bands,  where  the 
solutions  are  more  opaque.  Solutions  of  hydrated  salts  are  in  general 
more  transparent  than  pure  water.  All  things  considered,  we  regard 
this  as  the  strongest  evidence  thus  far  obtained  in  favor  of  the  solvate 
theory  of  solution. 

The  work  of  Dr.  Smith,  on  the  conductivity  and  dissociation  of  cer- 
tain organic  acids  in  water,  is  a  continuation  of  that  which  has  already 
appeared  in  publication  No.  170  of  the  Carnegie  Institution  of  Wash- 
ington. The  investigation  by  Dr.  Wightman  and  Mr.  Wiesel,  on  the 
conductivity  of  organic  acids  in  alcohol,  is  a  continuation  of  the  work 
which  has  been  in  progress  in  this  laboratory  for  more  than  ten  years 
on  the  conductivity  and  dissociation  of  electrolytes  in  water  as  a  sol- 
vent. (See  publication  No.  170  of  the  Carnegie  Institution  of  Wash- 
ington.) While  this  investigation  is  only  preliminary,  results  of  interest 
have  already  been  obtained. 

Dr.  Wightman,  Dr.  Davis,  and  Mr.  Holmes  made  a  very  exhaustive 
study  of  two  simple  salts  in  mixtures  of  alcohol  and  water,  solutions  in 
mixtures  of  these  solvents  showing  abnormal  properties. 

The  work  by  Dr.  Davis  and  Dr.  Hughes,  on  the  properties  of  solu- 
tions in  acetone,  was  taken  up  because  of  the  abnormal  behavior  of 
acetone  as  a  solvent. 

The  investigation  by  Dr.  Davis  and  Mr.  Putnam,  of  ternary  mix- 
tures of  glycerol,  acetone,  and  water,  had  in  mind  the  fact  that  glycerol 
has  a  very  high  viscosity,  water  intermediate  viscosity,  and  acetone  a 
very  low  viscosity.  The  viscosities  and  conductivities  of  solutions  in 
ternary  mixtures  of  these  solvents  were  studied.  A  general  discussion 
of  the  results  obtained,  bearing  on  the  solvate  theory  of  solution, 
seemed  desirable.  The  work  as  completed  was  published  in  a  large 


4  PREFACE. 

number  of  papers,  and  in  a  fairly  large  number  of  journals  in  America, 
Germany,  England,  and  Switzerland.  A  general  discussion  of  the 
results  thus  far  obtained  would  render  reference  to  the  work  more 
convenient.  The  last  chapter  of  this  monograph  gives  in  concise  form 
such  a  discussion  and  summary.  A  bibliography  of  the  papers  and 
monographs  already  published,  bearing  upon  this  theory,  will  make 
reference  to  the  literature  simpler. 

Finally,  it  gives  me  great  pleasure  to  thank  the  Carnegie  Institution 
of  Washington  for  the  generous  aid  with  which  they  have  supported 
these  investigations,  and  without  which  it  would  have  been  impossible 
to  do  much  of  this  work. 

HAERY  C.  JONES. 


CONTENTS. 

CHAPTER  I. 

THE  ABSORPTION  SPECTRA  OF   AQUEOUS  SOLUTIONS  OP   HTDRATED   AND  NON- 
HYDRATED  SALTS,  AS  STUDIED  BY  MEANS  OF  THE  RADIOMICROMETER. 

The  Radiomicrometer 9 

Method  of  Procedure 10 

The  Cells 11 

Earlier  Results 13 

Effect  of  Slit-Width 13 

Water  Absorption 16 

Effect  of  Slit-Width  on  the  Absorption  of  Light  by  Water 18 

Absorption  of  Light  by  Water  as  Affected  by  Hydrated  and  by  Non-hydrated  Salts .  19 

Method  of  Procedure 20 

Results 20 

Discussion  of  the  Results 33 

CHAPTER  II. 

CONDUCTIVITIES,  TEMPERATURE  COEFFICIENTS  OF  CONDUCTIVITY,  DISSOCIATIONS, 
AND  CONSTANTS  OF  CERTAIN  ORGANIC  ACIDS  IN  AQUEOUS  SOLUTIONS. 

Introduction 47 

Purpose  of  the  Investigation 48 

Experimental 48 

Reagents 49 

Apparatus 49 

Procedure 51 

Cell  Constants 51 

Values 54 

Limiting  Conductivities  of  the  Acids 55 

Malic  acid 57 

Aconitic  acid 57 

w-Chlorobenzqic  acid 57 

p-Chlorobenzoic  acid 57 

o-Bromobenzoic  acid 57 

m-Bromobenzoic  acid 57 

m-Acetoxybenzoic  acid 57 

o-Sulphobenzoic  acid 58 

m-Sulphobenzoic  acid 58 

p-Aminobenzensulphonic  acid 58 

Sebasic  acid 58 

Dissociation  Constants 59 

Summary 62 

CHAPTER  III. 

A  PRELIMINARY  STUDY  OF  THE  CONDUCTTVITY  OF  CERTAIN  ORGANIC  ACIDS  IN  ETHYL 
ALCOHOL  OF  15°,  25°,  AND  35°. 

Historical 64 

Experimental 67 

Results 73 

Malonic  acid 73 

o-Chlorobenzoic  acid 73 

p-Chlorobenzoic  acid 74 

p-Bromobenzoic  acid 74 

o-Nitrobenzoic  acid 74 

n^Titrobenzoic  acid 74 

,  4,  Dinitrobenzoic  acid 75 

1,  2,  4,  Dihydroxybenzoic  acid 75 

Tetrachlorphthalic  acid ; 75 

Malonic,  and  Change  in  Concentration 75 

o-Chlorobenzoic  and  Change  in  Concentration 76 

p-Chlorobenzoic  and  Change  in  Concentration 76 

p-Bromobenzoic  and  Change  in  Concentration 76 

5 


6  STUDY   OF   ABSORPTION   SPECTRA. 

Results^-Continued.  Pa*e 

o-Nitrobenzoic  and  Change  in  Concentration 76 

p-Nitrobenzoic  and  Change  in  Concentration 76 

1,  2,  4,  Dinitrobenzoic  and  Change  in  Concentration 76 

1,  2,  4,  Dihydroxybenzoic  and  Change  in  Concentration 76 

Tetrachlorphthalic  acid  and  Change  in  Concentration 76 

Discussion  of  Results 77 

CHAPTER  IV. 

THE  CONDUCTIVITY  AND  VISCOSITY  OF  SOLUTIONS  OF  POTASSIUM  IODIDE  AND  SODIUM 
IODIDE  IN  MIXTURES  OF  ETHYL  ALCOHOL  AND  WATER. 

Experimental 80 

Pure  Anhydrous  Alcohol 80 

Specific-Gravity  Determinations 80 

Mixed  Solvents 81 

Dissolved  Salts 82 

82 

82 

83 

84 

85 


Pipettes 

Conductivity  Cells 

Temperature  Regulators . 


Corrections  for  Expansion  and  Contraction 

Viscosities 

Viscosity  and  Fluidity  of  Potassium  Iodide  in  Alcohol- Water  Mixtures 86 

Temperature  Coefficients  of  Fluidity  of  Potassium  Iodide 87 

Viscosity  and  Fluidity  of  Sodium  Iodide  in  Alcohol-Water  Mixtures 87 

Conductivity  of  Potassium  Iodide  in  Mixtures  of  Ethyl  Alcohol  and  Water. ...  88 

Conductivity  of  Sodium  Iodide  in  Mixtures  of  Ethyl  Alcohol  and  Water 88 

Discussion  of  the  Results 89 

Viscosity  and  Fluidity 89 

Conductivity 91 

Summary 96 

Conclusions 96 

CHAPTER  V. 

CONDUCTIVITY  AND  VISCOSITY  OF  SOLUTIONS  OF  RUBIDIUM  SALTS  IN  MIXTURES  OF 
ACETONE  AND  WATER. 

Experimental 98 

Conductivity  Apparatus 98 

Bridge 98 

Cells 99 

Constant  Temperature  Baths 99 

Temperature  Regulation 99 

Viscosity  Apparatus 100 

Solvents 101 

Water 101 

Acetone 101 

Mixtures  of  Acetone  and  Water 101 

Solutions 

Salts 

Procedure. . . 


Conductivity  Measurements .  . 
Temperature  Coefficients . 


102 

102 

102 

102 

103 

Molecular  Conductivity  of  Rubidium  Chloride 104 

Molecular  Conductivity  of  Rubidium  Bromide 104 

Molecular  Conductivity  of  Rubidium  Iodide 105 

Molecular  Conductivity  of  Rubidium  Nitrate 106 

Viscosity  and  Fluidity  of  Rubidium  Chloride  and  Temperature  Coefficients 

of  Fluidity 107 

Viscosity  and  Fluidity  of  Rubidium  Bromide  and  Temperature  Coefficients 

of  Fluidity 108 

Viscosity  and  Fluidity  of  Rubidium  Iodide  and  Temperature  Coefficients 

of  Fluidity 109 

Viscosity  and  Fluidity  of  Rubidium  Nitrate  and  Temperature  Coefficients 

of  Fluidity 110 

Comparison  of  the  Viscosity  and  Fluidity  values  of  Rubidium  Bromide  in 

Acetone-water  Mixtures Ill 

Discussion Ill 

Summary 116 


CONTENTS.  7 

CHAPTER  VI. 

THE  CONDUCTIVITY  AND  VISCOSITY  OF  CERTAIN  RUBIDIUM  AND  AMMONIUM  SALTS  IN 
TERNARY  MIXTURES  OF  GLYCEROL,  ACETONE,  AND  WATER  AT  15°,  25°,  AND  35°. 

Page. 

Introduction 117 

Experimental 119 

Apparatus 119 

Thermostats 119 

Conductivity  Apparatus 120 

Viscosity  Apparatus ....". 120 

Volumetric  Apparatus 120 

Solvents 120 

Salts 122 

Solutions 122 

Procedure 122 


Viscosity  Data.  .  .  . 
Conductivity  Data. 


Discussion  of  Results . 


Summary 


123 
127 
131 
140 


CHAPTER  VII. 


DISCUSSION  OF  EVIDENCE  ON  THE  SOLVATE  THEORY  OF  SOLUTION  OBTAINED  IN  THE 
LABORATORIES  OF  THE  JOHNS  HOPKINS  UNIVERSITY. 

Earlier  Work 141 

Relation  between  Lowering  of  the  Freezing-point  of  Water  and  Water  of  Crystalli- 
zation of  the  Dissolved  Substance 143 

Approximate  Composition  of  the  Hydrates  formed  by  Various  Substances  in  Solution  145 
Relation  between  the  Minima  in  the  Freezing-point  Curves  and  the  Minima  in  the 

Boiling-point  Curves 147 

Relation  between  Water  of  Crystallization  and  Temperature  of  Crystallization. . . .  148 
Hydrate  Theory  in  Aqueous  Solutions  becomes  the  Solvate  Theory  in  Solutions  in 

General 149 

Temperature  Coefficients  of  Conductivity  and  Hydration 151 

Relation  between  the  Hydration  of  the  Ions  and  Their  Ionic  Volumes 155 

Hydration  of  the  Ions  and  the  Velocities  with  which  They  Move 157 

Dissociation  as  Measured  by  the  Freezing-point  Method  and  by  the  Conductivity 

Method .    158 

Effect  of  one  Salt  with  Hydrating  Power  on  the  Hydrates  formed  by  a  Second  Salt 

in  the  Same  Solution 161 

Investigations  in  Mixed  Solvents 162 

Spectroscopic  Evidence  Bearing  on  the  Solvate  Theory  of  Solution 170 

Work  of  Jones  and  Uhler 170 

Work  of  Jones  and  Anderson 172 

Work  of  Jones  and  Strong 175 

Absorption  Spectra  of  Neodymium  Salts 178 

a  Bands 180 

Effect  of  Rise  in  Temperature 181 

Spectrophotography  of  Chemical  Reactions 183 

Work  of  Jones  and  Guy  on  the  Absorption  Spectra  of  Solutions 185 

Work  of  Jones,  Shaeffer,  and  Paulus 188 

Summary  of  the   Lines  of  Evidence  obtained  in  this  Laboratory  bearing  on  the 

Solvate  Theory  of  Solution 189 

How  the  present  Solvate  Theory  of  Solution  differs  from  the  older  Hydrate  Theory . .  190 

Significance  of  the  Solvate  Theory  of  Solution 192 

The  Solvate  Theory  and  the  Theory  of  Electrolytic  Dissociation 193 

Does  the  Solvate  Theory  help  to  explain  any  of  the  Apparent  Exceptions  to  the 

Theory  of  Electrolytic  Dissociation? 194 

Does  the  Solvate  Theory  aid  us  in  Explaining  the  Facts  of  Chemistry  in  General 

and  Physical  Chemistry  in  Particular? 196 

Why  is  the  Nature  of  Solutions  of  such  Vital  Importance  not  only  for  Chemistry 

but  for  Science  in  General? 197 

Bibliography 199 

Papers 199 

Monographs 202 


CHAPTER  I. 

THE  ABSORPTION  SPECTRA  OF  AQUEOUS  SOLUTIONS  OF  HYDRATED 

AND  NON-HYDRATED  SALTS,  AS  STUDIED  BY  MEANS 

OF  THE  RADIOMICROMETER. 

BY  E.  J.  SHAEFFER  AND  M.  G.  PATTLUS. 

The  object  in  using  the  radiomicrometer  for  studying  the  absorption 
spectra  of  solutions  was  to  work  quantitatively.  The  photographic 
method  gave  us  the  positions  of  the  various  bands,  and  a  qualitative,  or, 
at  best,  very  roughly  quantitative  estimate  of  the  relative  intensities 
of  the  various  lines  and  bands.  The  personal  equation  of  the  photo- 
graphic plate,  however,  always  comes  into  play  and,  all  things  con- 
sidered, we  regard  the  photographic  method  as  useful  and  reliable 
mainly  in  determining  the  positions  of  the  various  lines  and  bands. 

A  more  important  problem  and  one  more  fundamental  than  the 
determination  of  the  positions  of  the  lines  and  bands,  is  the  measure- 
ment of  the  relative  intensities  of  these  lines  and  bands  and  the  relative 
intensities  of  the  different  parts  of  the  same  band.  For  this  purpose 
some  instrument  must  be  used  which  measures  the  intensity  of  the 
radiation  falling  upon  it.  We  selected  the  radiomicrometer  as  best 
adapted  to  this  purpose,  it  being  a  thermo-electric  junction  attached 
to  a  loop  of  non-magnetic  wire,  and  the  whole  suspended  in  a  magnetic 
field.  The  following  form  was  constructed  and  used: 

THE  RADIOMICROMETER; 

The  radiomicrometer  used  in  this  investigation  was  built  by  Dr.  Guy. 
It  had  a  full  period  of  8  seconds  and  a  sensibility  of  8  cm.  per  square 
millimeter  of  exposed  junction,  candle  and  scale  being  at  a  distance  of 
1  meter.  This  instrument,  therefore,  combines  very  high  sensibility 
with  very  short  period.  It  had  no  compensating  junction,  and  this 
necessitated  very  careful  protection  from  changes  in  temperature,  from 
air  drafts,  and  from  energy  radiations  in  general;  otherwise  the  zero- 
point  would  be  unsteady.  It  was  found  that  by  incasing  the  instru- 
ment in  a  tight  wooden  box  of  such  dimensions  that  the  radiomicrometer 
was  surrounded  by  a  layer  of  cotton  a  foot  thick,  the  drift  of  the  zero- 
point  was  entirely  eliminated  if  the  temperature  of  the  room  was  f  airly 
constant.  It  was  quite  practicable  to  keep  the  room  at  a  sufficiently 
constant  temperature  for  this  purpose.  The  accuracy  of  the  results 
was  greatly  increased  by  keeping  the  zero-point  constant.  Precaution 
was  also  taken  to  keep  the  current  which  passed  through  the  Nernst 
glower  as  invariable  as  possible.  The  source  of  electricity  used  was  a 
large  storage  battery,  which  supplied  a  very  constant  current,  provided 


10  STUDY    OF   ABSORPTION   SPECTRA 

the  Nernst  glower  was  carefully  protected  from  air-currents.  This 
was  accomplished  by  inclosing  the  glower  in  a  tight  box  of  heavy 
asbestos  board.  That  the  current  must  have  remained  invariable  during 
the  two  successive  readings  which  we  had  to  make,  to  obtain  the  per- 
centage transmission  for  any  given  layer  of  solution,  is  shown  by  the 
close  agreement  of  the  duplicate  readings.  As  a  rule,  the  duplicate 
results  agreed  with  one  another  to  within  less  than  1  per  cent.  That 
we  could  duplicate  our  readings  so  closely  we  regard  as  due  to  the  con- 
stancy of  the  zero-point  of  the  radiomicrometer  and  to  the  steadiness 
of  the  current  passing  through  the  Nernst  glower.  Another  factor 
which  seriously  interferes  with  the  accuracy  of  this  work  is  mechanical 
vibrations.  This  effect  becomes  especially  noticeable  when  the  deflec- 
tions of  the  radiomicrometer  are  small.  By  placing  felt  under  the 
marble  slab  which  served  as  the  base  of  support  of  the  instrument,  and 
setting  the  radiomicrometer  support  on  blocks  of  wood  and  cork,  we 
were  able  to  eliminate  all  mechanical  disturbances  except  those  of  a 
very  violent  character.  When  such  violent  disturbances  occurred,  the 
work  was  temporarily  discontinued  and  the  measurements  repeated. 

METHOD  OF  PROCEDURE. 

The  Nernst  glower  carried  0.8  ampere  and  120  volts.  By  means  of 
an  adjustable  resistance  we  could  keep  the  current  very  close  to  0.8 
ampere  throughout  this  entire  work.  If  any  appreciable  variation 
in  the  current  was  detected  between  the  two  successive  readings  which 
determined  the  transmission  for  any  given  layer  of  solution,  the  readings 
were  always  repeated.  The  light  from  the  Nernst  glower,  after  being 
rendered  parallel  by  passing  through  a  lens,  was  passed  through  a  given 
layer  of  solution  and  then  focused  on  the  slit  of  the  Hilger  spectroscope. 

By  properly  rotating  the  drum-head  of  the  spectroscope,  we  could 
throw  the  desired  wave-lengths  of  light  on  the  junction  of  the  radio- 
micrometer.  The  dispersion  curve  for  the  glass  prism  was  carefully 
worked  out,  some  of  the  points  in  the  infra-red  being  determined  by 
means  of  the  water-bands.  It  was  possible  to  find  definite  maxima  of 
absorption  for  the  water-bands  1/x,  1.25^,  LSI/*,  and  2.01/i.  Since  we 
worked  mainly  with  aqueous  solutions,  after  each  series  of  measure- 
ments we  could  determine,  by  means  of  the  water-bands,  whether  the 
prism  had  shifted  or  not. 

Suppose  it  is  desired  to  know  the  absolute  transmission  of  a  given 
solution,  20  mm.  in  depth,  for  a  given  wave-length  of  light.  This  was 
done  by  the  following  differential  method,  the  correctness  of  which  has 
already  been  discussed:1  Twenty-one  millimeters  of  the  solution  were 
placed  in  one  cell  and  1  mm.  was  placed  in  a  second  cell,  which  was 
exactly  like  the  first  one.  The  wave-length  of  light,  whose  transmis- 

'Carnegid  Inst.  Wash.  Pub.  190,  p.  62. 


BY   MEANS   OF  THE   RADIOMICROMETER.  11 

sion  by  the  solution  it  was  desired  to  measure,  was  made  to  fall  upon 
the  junction  after  passing  through  the  solution,  by  turning  the  cali- 
brated drum-head  to  the  proper  point.  The  deflection  when  the  deeper 
layer  was  in  the  path  of  the  beam  of  light  was  noted,  and  as  quickly  as 
possible  the  cell  containing  the  more  shallow  layer  was  placed  in  the 
position  formerly  occupied  by  the  deeper  layer.  This  was  accom- 
plished by  a  carefully  adjusted  sliding  carriage.  The  deflection  when 
the  more  shallow  layer  was  in  the  path  of  the  beam  of  light  was  noted. 
The  deflection  produced  when  the  light  passed  through  the  deeper 
layer  of  solution,  divided  by  the  deflection  when  the  more  shallow  layer 
was  in  the  path  of  the  light,  gave  the  absolute  percentage  transmission 
of  20  mm.  of  the  solution,  for  the  wave-length  of  light  in  question. 

In  work  such  as  this,  involving  as  it  does  so  many  distinct  and 
separate  operations,  it  is  highly  desirable  to  duplicate  all  of  the  meas- 
urements. This  has  been  done  in  nearly  all  of  the  work  the  results 
of  which  are  herein  recorded.  In  only  some  cases  are  the  duplicate 
results  given.  In  every  series  of  measurements,  duplications  were 
made  for  certain  wave-lengths  of  light.  After  filling  the  cells  with 
solution  or  solvent,  and  making  the  necessary  adjustments  as  accu- 
rately as  possible,  the  percentage  transmission  was  determined  several 
tunes  for  wave-length  X  706.  When  20  mm.  of  the  solution  or  solvent 
was  employed  this  was  always  about  95.5  per  cent.  After  taking  about 
ten  readings  farther  down  in  the  infra-red,  wave-length  of  light  X  706 
was  again  passed  through  the  solution  and  thrown  upon  the  junction. 
The  transmission  for  this  wave-length  was  repeatedly  determined  dur- 
ing the  entire  series  of  measurements.  Usually  we  obtained  the  same 
percentage  transmission  in  the  two  cases,  to  within  the  limit  of  error  of 
the  method.  This  showed  that  comparable  conditions  were  maintained 
in  both  cells  during  the  time  required  to  make  the  intermediate  read- 
ings. Duplicate  results  are  given  for  the  wave-lengths  X=0.706ju, 
X=1.00ju,  and  X  =  1.24jU.  These  were  secured  after  the  entire  series  of 
measurements  had  been  completed.  To  obtain  reliable  readings  great 
care  had  to  be  exercised  in  adjusting  the  cells  in  the  carriage,  so  that 
there  was  the  same  distribution  of  light  from  the  glower  upon  the 
slit  when  either  cell  was  placed  in  the  path  of  the  light. 

THE  CELLS. 

The  cells  for  holding  the  solutions  and  solvent  are  among  the  most 
important  parts  of  the  apparatus,  and  the  reliability  of  the  results  is 
largely  dependent  upon  how  they  meet  various  requirements.  Two 
cells  were  made  for  this  work,  and  they  were  made  as  nearly  alike  as 
possible.  Each  cell  consisted  essentially  of  two  brass  cylinders,  which 
telescoped  neatly  into  one  another.  Glass  plates  were  set  into  the 
ends  of  each  cylinder  by  means  of  Wood's  fusible  metal,  so  that  then* 
surfaces  were  parallel  to  one  another.  The  question  as  to  whether  the 


12 


STUDY   OF   ABSORPTION    SPECTRA 


plates  were  sufficiently  piano-parallel  was  determined  by  the  character 
of  the  interference  fringes  which  they  gave.  These  glass  plates  closing 
the  ends  of  the  cells  were  1  mm.  thick.  We  found  some  difficulty  in 
procuring  plates  of  the  best  optical  glass  which  were  piano-parallel. 
Many  times  when  the  plates  were  sufficiently  piano-parallel  for  our 
purpose,  the  Wood's  metal  surrounding  the  edges  of  the  plates  would, 
on  cooling  or  setting,  so  warp  the  plate  that  it  was  necessary  to  reset  it. 
This  often  had  to  be  repeated  several  times  to  secure  the  desired  result. 

The  distance  between  the  glass  ends  was  regulated  by  means  of  a 
finely  threaded  nut  which  screwed  on  to  the  outer  cylinder.  Each 
complete  revolution  of  this  nut  raised  or  lowered  the  inner  cylinder 
just  1  mm.  The  nut  was  calibrated  in  100  divisions  by  means  of  a 
dividing  engine,  and  with  this  nut  we  could  readily  and  accurately 
adjust  the  distance  between  the  glass  ends  to  less  than  0.01  mm. 

The  brass  cells  were  first  heavily  plated  with  silver,  being  taken  out 
of  the  plating  bath  from  time  to  time  and  thoroughly  rubbed  with  the 
finest  crocus  powder.  A  heavy  gold  plate  was  then  deposited  on  this 
silver  surface.  Unless  the  gold  was  of  sufficient  thickness  to  cover 
well  all  of  the  exposed  portions  of  the  Wood's  metal,  it  was  found  that 
the  concentrated  and  strongly  hydrolyzed  solutions  which  we  studied 
would  act  upon  the  Wood's  metal,  giving  rise  to  streamers  in  the 
solutions  and  often  to  an  opalescence.  This  would,  of  course,  lead  to 
considerable  error.  The  solutions  used  in  making  the  measurements 
were  always  perfectly  clear,  and  if  any  opalescence  developed  in  them 
while  making  the  measurements,  the  solutions  were  discarded  and  the 
cells  were  at  once  replated. 

TABLE  I.— Testing  the  two  cells. 


X 

Cell  I. 

Cell  II. 

X 

Cell  I. 

Cell  II. 

706 

27.5 

27.5 

1066 

123.0 

123.0 

746 

38.0 

38.0 

1138 

135.5 

136.0 

787 

50.5 

50.5 

1216 

63.5 

63.5 

833 

65.0 

65.0 

1292 

65.0 

65.0 

886 

79.5 

79.0 

1362 

45.5 

45.5 

941 

91.5 

92.0 

1429 

7.0 

7.0 

1003 

79.5 

79.5 

The  two  cells  were  tested  from  time  to  tune  to  see  if,  under  com- 
parable conditions,  they  both  gave  the  same  deflections.  This  was 
usually  done  by  placing  10  mm.  of  water  in  each  cell  and  noting  the 
actual  deflections  as  given  by  the  radiomicrometer  for  the  various 
wave-lengths  of  light.  Table  1  contains  the  results  for  one  such  test. 
The  given  wave-lengths  of  light  would  be  obtained  by  turning  the 
drum-head  of  the  spectrometer  at  intervals  of  ten  divisions  over  the 
entire  region  of  the  spectrum  which  we  studied.  In  all  tests  the 
deflections  for  the  two  cells  agreed  as  closely  as  those  given  in  table  1. 


BY   MEANS    OF   THE    RADIOMICROMETER.  13 

If  equal  depths  of  a  concentrated  aqueous  solution  were  used  in 
place  of  equal  depths  of  pure  water,  a  number  of  precautions  were 
necessary  in  order  that  the  same  deflections  would  be  obtained  with 
the  two  cells.  It  was,  of  course,  necessary  that  the  solutions  in  both 
cells  should  remain  perfectly  clear.  If  even  the  slightest  opalescence 
developed  in  one  of  the  cells,  the  transmission  in  this  cell  would  be 
less  than  in  the  other.  It  was  further  found  that  special  precautions 
had  to  be  taken  to  keep  the  glass  ends  of  the  cells  clean.  The  ordi- 
nary methods  of  cleaning  did  not  suffice.  In  filling  the  cells  with  a 
solution,  some  of  the  solution  will  usually  come  in  contact  with  the 
outer  surface  of  the  glass  plate  in  the  larger  cylinder.  Even  after 
considerable  rubbing  with  a  cleansing-cloth  there  will  remain  an  almost 
invisible  film  of  the  crystallized  salt  on  the  plate.  This  was  found  to  be 
sufficient  to  change  the  transmission  as  much  as  5  per  cent.  Special 
precautions  must  be  taken  to  clean  the  ends,  otherwise  comparable 
results  could  not  be  obtained. 

EARLIER  RESULTS. 

It  was  earlier  found  by  Jones  and  Guy1  that  when  the  product  of 
the  concentration  of  the  solution  multiplied  by  the  depth  of  layer  was 
kept  constant,  Beer's  law  did  not  hold  for  neodymium  salts.  Certain 
tentative  suggestions  were  made  in  that  work  which,  in  the  light  of  the 
present  results,  must  be  somewhat  modified.  The  point  in  question 
involves  the  increase  in  the  intensity  of  the  neodymium  bands  with 
dilution.  A  possible  explanation  of  this  phenomenon,  based  upon 
resonance,  was  offered. 

It  was  later  found  that  there  should  be  some  correction  made  for 
the  water  absorption  in  the  case  of  the  0.87/x  band,  since  here  the 
1/x  water-band  began  to  absorb  slightly.  If  we  take  into  account  both 
of  these  factors,  together  with  the  additional  correction  for  the  width 
of  slit2  that  was  used,  it  is  thought  that  the  phenomenon  referred  to 
above  might  be  accounted  for,  and  that  Beer's  law  would  hold  for  the 
dilute  solutions  of  neodymium  salts. 

EFFECT  OF  SLIT-WIDTH. 

The  effect  of  the  width  of  the  slit  on  the  character  of  the  transmission 
curves  for  solutions  of  neodymium  salts  was  not  well  understood.  It 
seemed  desirable  that  some  work  should  be  done  on  this  problem,  with 
the  hope  that  we  might  arrive  at  a  better  understanding  of  this  effect. 
It  was  found  to  be  especially  difficult  to  obtain  concordant  results 
when  working  with  neodymium  salts.  This  was  due  to  the  low  dis- 
persive power  of  the  glass  prism  with  which  the  Hilger  spectroscope 
was  provided,  to  the  difficulty  in  accurately  setting  the  calibrated  head, 
and  to  the  very  narrow  and  sharp  bands  of  neodymium  salts.  Rotating 

iAmer.  Chem.  Journ.,  44,  1913.  "Carnegie  Inst.  Wash.  Pub.  No.  190,  p.  70. 


14  STUDY   OF   ABSORPTION   SPECTRA 

the  calibrated  drum-head  one  division  towards  the  red  would  increase 
the  wave-length  of  light  about  40  Angstrom  units  for  the  region  which 
we  studied.  Even  this  slight  increase  in  wave-length  would  often 
change  the  percentage  transmission  as  much  as  10  to  20  per  cent.  The 
tabular  data  bring  this  out  very  clearly.  It  is  therefore  necessary  to 
exercise  great  care  in  setting  the  calibrated  head  which  rotates  the  glass 
prism.  Table  2  contains  data  which  show  the  effect  of  increasing 
the  slit-width,  the  changes  that  take  place  when  the  conditions  for 
Beer's  law  are  fulfilled,  and  the  accuracy  which  is  possible  in  setting 
the  calibrated  head  of  the  spectroscope  and  in  reading  the  deflections. 
We  could  not  read  the  deflections  closer  than  0.5  mm.,  and  in  the 
visible  red,  with  0.2  mm.  slit-width,  the  deflections  were  so  small  that 
the  error  in  reading  may  be  as  much  as  2  per  cent.  For  conditions 
such  as  are  indicated  by  this  table,  the  minimum  deflection  obtained 
for  a  slit-width  of  0.2  mm.  was  12  mm.;  that  for  0.4  mm.  slit-width 
was  63  mm.;  the  maximum  deflection  for  0.2  mm.  slit- width  was 
69  mm. ;  that  for  0.4  mm.  slit-width  was  378  mm. 

We  desired  to  use  slit-widths  as  large  as  possible  so  as  to  increase 
the  accuracy  of  the  readings ;  this,  however,  could  be  done  only  when 
the  wider  slit-widths  had  but  little  effect  on  the  percentage  transmission. 

It  is  desired  to  call  attention  to  the  duplicate  results,  which,  in 
the  tables,  are  bracketed.  For  each  duplicate  measurement  the  head 
of  the  spectrometer  was  reset  and  the  deflection  again  read.  The 
second  reading  was  made  after  the  series  of  readings  had  been  com- 
pleted, and  then,  after  completing  the  second  series  of  readings  a  third 
reading  was  taken  for  a  few  wave-lengths  of  light.  The  method  of 
procedure  was  as  follows :  It  was  desired  to  measure  the  absolute  trans- 
mission of  2  mm.  of  a  0.586  normal  solution  of  neodymium  chloride  in 
water.  This  was  done  by  the  differential  method  already  discussed. 
We  place  in  cell  1, 3  mm.  of  the  solution  and  note  the  deflection  as  given 
by  the  radiomicrometer.  Similarly,  we  place  1  mm.  of  the  solution  in 
cell  II  and  note  the  deflection  when  cell  II  occupies  exactly  the  same 
position  formerly  occupied  by  cell  I.  The  deflection  given  by  3  mm.  of 
the  solution,  divided  by  the  deflection  given  by  1  mm.  of  the  solution, 
gives  the  absolute  percentage  transmission  of  2  mm.  of  the  solution. 

The  solution  was  diluted  eight  times,  its  concentration  then  being 
0.073  normal.  To  satisfy  the  requirements  of  Beer's  law  we  would 
have  to  place  in  one  cell  24  mm.  of  the  more  dilute  solution  and  in  the 
other  cell  8  mm.  We  would  then  have  the  same  number  of  absorbers 
in  the  path  of  the  beam  of  light  as  we  had  in  the  case  of  the  original, 
more  concentrated  solution.  In  the  case  of  the  more  dilute  solution 
we  would  obtain  the  absolute  transmission  for  16  mm.  depth  of  layer. 
This  would  be  comparable  with  the  transmission  obtained  for  2  mm. 
of  the  more  concentrated  solution.  Column  1  in  table  2  gives  the 
actual  head-readings.  In  columns  2  and  3  are  given  the  percentage 


BY   MEANS   OF   THE   RADIOMICROMETER. 


15 


transmissions  for  the  more  concentrated  and  the  more  dilute  solutions, 
the  slit-width  being  0.2  mm.  Columns  4  and  5  contain  the  results  for 
the  two  solutions,  the  width  of  the  slit  being  0.3  mm.  Columns  6  and 
7  give  the  results  for  the  two  solutions  when  the  slit-width  was  0.4  mm. 
In  all  of  this  work  about  0.8  ampere  of  current  was  used. 

TABLE  2. — Results  with  different  widths  of  slil. 


1 

Head- 
read- 
ings. 

Slit-width  0.2  mm. 

Slit-width  0.3  mm. 

Slit-width  0.4  mm. 

2 

0.586  N. 

3 

0.073  N. 

4 

0.586  N. 

5 

0.073-  N. 

6 

0.586  N. 

7 
0.073  N. 

478 

f  100.0 

96.0 

100.0 

96.3 

100.0 

96.3 

478 

,  100.0 

94.0 

100.0 

95.5 

100.0 

95.6 

486 

78.0 

69.6 

80.1 

72.3 

80.3 

75.4 

487 

57.0 

55.8 

64.3 

60.3 

68.0 

64.5 

488 

47.0 

46.8 

51.3 

48.8 

56.8 

57.0 

489 

43.5 

40.4 

44.1 

41.0 

49.4 

45.7 

489 

44.4 

39.5 

44.4 

41.3 

49.6 

46.5 

489 

44.4 

40.4 

44.7 

41.5 

•  49.3 

46.4 

490 

'     49.0 

48.0 

50.0 

45.5 

50.6 

47.0 

490 

\     48.0 

48.0 

50.0 

45.5 

50.8 

47.4 

491 

63.3 

60.0 

58.6 

54.4 

58.7 

55.7 

492 

76.6 

72.5 

73.4 

66.7 

68.8 

64.5 

494 

94.7 

88.0 

92.9 

81.5 

86.6 

81.6 

500 

75.6 

62.2 

77.7 

75.0 

77.7 

72.4 

501 

58.6 

54.7 

65.7 

59.5 

68.7 

62.6 

502 

/     46.3 

43.5 

53.3 

47.7 

59.4 

54.5 

502 

1     47.0 

43.5 

53.4 

48.3 

60.0 

54.1 

503 

45.4 

43.0 

50.0 

44.5 

54.3 

49.7 

503 

44.8 

43.0 

49.7 

44.7 

54.4 

49.7 

503 

44.2 

43.0 

49.4 

45.0 

54.4 

49.3 

504 

f     58.0 

53.0 

55.5 

50.5 

55.8 

52.0 

504 

1     57.3 

54.2 

55.4 

51.0 

56.0 

52.3 

506 

79.0 

74.0 

77.6 

70.0 

71.7 

67.3 

517 

77.3 

69.4 

80.8 

70.7 

81.1 

73.0 

517 

76.5 

69.0 

81.6 

71.0 

81.5 

72.5 

518 

76.6 

68.0 

78.6 

68.6 

78.9 

70.3 

518 

75.9 

68.3 

78.8 

68.8 

78.8 

70.6 

518 

76.4 

68.7 

79.3 

67.5 

69.7 

519 

78.3 

70.3 

79.7 

68.4 

78!6 

519 

78.3 

70.7 

79.9 

69.0 

78.3 

519 

79.2 

69.7 

79.6 

69.3 

7613 

520 

82.4 

74.0 

81.7 

71.1 

79.3 

71.0 

From  the  transmission  values  given  in  table  3,  we  can  gain  a  fairly 
accurate  conception  of  the  intensity  of  the  absorption  due  to  the  water 
contained  in  2  mm.  of  the  more  concentrated  solution,  and  in  16  mm. 
of  the  more  dilute  solution.  The  water  absorption  was  taken  only  for 
those  wave-lengths  which  correspond  to  the  centers  of  the  neodymium 
absorption  bands. 

Tables  2  and  3  bring  out  several  interesting  points  concerning  the 
neodymium  bands.  The  effect  of  increasing  the  slit-width  is  to  dis- 
place the  bands  towards  the  longer  wave-lengths.1  As  we  approach 

^his  is  due  to  the  slit  having  but  one  movable  edge,  the  other  being  stationary. 


16 


STUDY   OF   ABSORPTION   SPECTRA 


the  bands  the  transmission  is  greatest  with  the  widest  slit.  As  we 
leave  the  bands  the  transmission  is  least  with  the  widest  slit.  This  is 
to  be  expected,  since  the  spectrum  lines  have  their  centers  displaced 
slightly  towards  the  red  with  increasing  slit-width.  It  will  be  readily 
seen  that  the  neodymium  bands  are  exceedingly  sharp  and  narrow. 
Doubling  the  slit-width  will,  for  some  wave-lengths  of  light,  change 
the  transmission  15  per  cent  or  more.  It  is  clear  that  the  centers  of 
the  bands  do  not  become  more  intense  with  increasing  slit-widths. 

TABLE  3.— Water  absorption. 


Head- 
read- 
ings. 

1.92mm., 
H20. 

15.89  mm., 
H20. 

Head- 
read- 
ings. 

1.92  mm., 
H20. 

15.89  mm., 
H20. 

489 

f     99.0 

96.6 

504 

98.0 

98.0 

489 

99.0 

95.8 

518 

f     98.0 

93.1 

489 

99.0 

96.6 

518 

\     98.0 

93.3 

503 

|     99.0 

98.0 

518 

[     98.0 

93.3 

503 

99.0 

98.0 

520 

97.0 

92.5 

503 

1     99.0 

98.5 

Indeed,  just  the  opposite  effect  is  noted.  Table  2  also  shows  that, 
aside  from  considerations  of  the  purity  of  the  spectrum,  the  slits  of 
the  spectroscope  we  used  should  actually  be  of  almost  infinitesimal 
width  for  lines  and  bands  as  sharp  and  narrow  as  those  of  neodymium. 
If  the  slit  was  narrower  than  0.2  mm.,  the  resulting  deflections  would 
be  so  small  that  the  error  in  reading  the  scale  would  be  considerable. 
The  effect  of  the  width  of  the  slit  on  the  absorption  bands  of  water 
was  also  studied,  and  will  be  discussed  later. 

WATER  ABSORPTION. 

If  we  consider  the  water  absorption  we  find  that  the  neodymium 
bands  increase  in  intensity  only  slightly  with  dilution.  Take,  for 
example,  the  band  represented  by  head-reading  489;  2  mm.  of  solution 
have  less  than  1  per  cent  absorption  at  this  point,  due  to  the  water 
present.  It  was  found  from  the  specific  gravity  and  concentration  of 
the  solution  that  95.6  per  cent  by  volume  was  water.  Therefore,  one 
cell  was  filled  with  2.87  mm.  of  water  and  the  second  with  0.95  mm. 
of  water.  Each  cell,  then,  had  just  as  much  water  in  it  as  it  had  when 
filled  with  3  and  1  mm.,  respectively,  of  the  solution.  In  this  way  the 
absolute  transmission  for  a  layer  of  water  equal  to  that  in  the  2  mm. 
of  solution  is  determined. 

Considering  now  the  more  dilute  solution,  the  cell  which  contained 
24  mm.  of  solution  actually  contained  23.83  mm.  of  water,  the  percent- 
age of  water  present  being  99.3.  The  second  cell  contained  7.94  mm. 
of  water.  By  the  differential  method  we  thus  find  the  transmission 
for  a  layer  of  water  equal  to  that  contained  in  16  mm.  of  the  more 
dilute  solution.  The  particular  region  which  we  are  considering  is  in 


BY   MEANS   OF   THE    RADIOMICROMETER. 


17 


the  neighborhood  of  the  0.77ju  water-band.  This  is  a  very  faint  band, 
and  even  for  the  above  depth  of  water  the  actual  absorption  is  only 
between  3  and  4  per  cent.  It  will  thus  be  seen  that  for  the  first  band 
whose  center  is  represented  by  head-reading  489,  the  concentrated  and 
dilute  solutions  absorb  with  equal  intensity. 

Wave-length  of  light  represented  by  head-reading  503  is  the  point 
which  indicates  the  center  of  the  first  neodymium  band  in  the  infra- 
red. With  slit-width  of  0.2  mm.,  the  water  contained  in  2  mm.  of  the 
more  concentrated  solution  has  an  absorption  of  1  per  cent  for  this 
particular  wave-length,  while  the  water  contained  in  16  mm.  of  the 
more  dilute  solution  has  an  absorption  of  2  per  cent.  Here,  as  in  the 
case  of  the  other  bands,  the  intensity  of  the  bands  for  the  dilute  and 
concentrated  solutions  is  the  same. 

TABLE  4. — Testing  Beer's  Law. 


X 

Cone.  0.6 
normal, 

depth  2  mm. 

Cone.  0.75 
normal, 
depth  16  mm. 

X 

Cpnc.  0.6 
normal, 

depth  2  mm. 

Cone.  0.75 
normal, 
depth  16  mm. 

702 

100 

100 

803 

/    39.3 

39.8 

702 

100 

100 

803 

\    39.3 

39.3 

732 

63.5 

61.3 

807 

/    59.5 

56.5 

736 

55.0 

52.5 

807 

\    60.5 

56.3 

739 

/    41.6 

39.0 

813 

76.3 

72.0 

739 

i    41.6 

38.0 

817 

84.0 

80.0 

743 

/    40.5 

37.0 

869 

/    84.0 

77.0 

743 

\    40.5 

36.8 

869 

1    83.5 

76.3 

749 

44.5 

875 

I    73.5 

65.5 

753 

66.5 

64.2 

875 

\    73.8 

66.0 

756 

82.0 

74.5 

880 

f    76.5 

68.5 

764 

94.3 

90.0 

880 

\    77.0 

69.0 

789 

73.5 

66.7 

885 

86.0 

77.3 

794 

42.3 

42.3 

890 

93.0 

81.7 

799 

f    34.7 

31.2 

799 

1    34.7 

32.5 

Head-reading  518  marks  the  position  for  the  last  neodymium  band 
we  could  study.  The  wave-length  of  light  corresponding  to  this  band 
would  be  approximately  0.87/z,  which  is  sufficiently  close  to  the  I/* 
water-band  for  the  dilute  solution  to  cause  considerable  absorption  on 
account  of  the  water  present.  For  this  wave-length  16  mm.  of  the 
more  dilute  solution  has  about  7  per  cent  absorption  due  to  water; 
whereas  the  2  mm.  of  the  more  concentrated  solution  has  2  per  cent 
absorption  at  this  wave-length.  It  will  thus  be  seen  that  only  in  the 
case  of  this  one  band  is  there  an  increase  in  intensity  with  dilution. 
The  more  dilute  solution  has  a  band  which  is  nearly  3  per  cent  more 
intense  than  that  for  the  more  concentrated  solution. 

It  should  also  be  noted  that  the  absorption  bands  of  the  more  dilute 
solution  are  broader  than  those  of  the  more  concentrated.  This  is 
shown  by  the  fact  that  the  percentage  transmission  of  the  more  dilute 
solutions  are  always  lower  than  those  of  the  more  concentrated.  There 
is  no  evidence  of  a  shift  in  the  bands  for  the  dilute  solution. 


18 


STUDY   OF   ABSORPTION    SPECTRA 


Owing  to  the  large  number  of  possible  sources  of  error,  it  was  regarded 
as  desirable  to  repeat  the  work  bearing  on  the  point  under  discussion. 
This  was  done  three  times;  and  that  the  results  obtained  are  very  con- 
cordant can  be  seen  by  comparing  table  4  with  table  3.  The  slit-width 
used  was  0.2  mm.  The  depths  of  layers  were  the  same  as  in  table  3, 
and  the  concentration  of  the  new  solution  was  six-tenths  normal. 

The  results  seem  to  justify  the  conclusion  that  the  intensities  of  the 
bands  for  both  the  dilute  and  concentrated  solutions  are  practically  the 
same,  and  that  Beer's  law  holds  fairly  well  for  solutions  of  neodymium 
chloride.  The  bands  in  the  case  of  the  more  dilute  solution  are  broader 
and  only  very  slightly  more  intense  than  for  the  more  concentrated. 

EFFECT  OF  SLIT-WIDTH  ON  THE  ABSORPTION  OF  LIGHT  BY  WATER. 

It  will  be  remembered  that  increasing  the  width  of  the  slit  greatly 
changed  the  transmission  of  salts  of  neodymium.  Since  we  were  mak- 
ing an  extensive  study  of  the  water-bands,  it  was  desired  to  know  what 
changes  would  take  place  in  the  transmission  of  water  with  increasing 
slit-widths.  The  In  water-band  compared  with  the  neodymium  bands 
is  relatively  broad,  and  the  1.25^  water-band  is  very  broad,  it  being 
difficult  to  find  accurately  a  maximum  of  absorption.  To  study  the 
1/z  water-band  a  depth  of  20  mm.  was  employed,  while  10  mm.  sufficed 
to  study  the  1.25/z  band  of  water. 

TABLE  5.— Effect  of  slit-width. 


IM  water-band. 

l.25n  water-band. 

Slit-width 

ShVwddth 

Slit-width 

Slit-width 

Slit-width 

Slit-width 

0.2  mm. 

0.3  mm. 

0.4  mm. 

0.2  mm. 

0.3  mm. 

0.4  mm. 

941 

71.5 

72.6 

74.8 

1066 

73.5 

73.5 

72.9 

952 

63.5 

62.8 

62.4 

1138 

67.5 

70.0 

69.0 

965 

51.6 

52.7 

57.6 

1216 

24.7 

26.0 

28.3 

978 

41.6 

42.5 

47.0 

1231 

23.5 

23.2 

24.9 

984 

38.1 

38.8 

44.0 

1238 

23.2 

23.0 

24.0 

990 

35.6 

36.6 

40.3 

1246 

23.1 

22.9 

23.8 

996 

34.6 

35.1 

38.8 

1254 

23.1 

22.9 

23.2 

1003 

34.8 

35.4 

37.4 

1261 

23.2 

22.8 

23.1 

1010 

35.5 

36.1 

37.6 

1277 

23.4 

23.6 

23.8 

1017 

37.4 

37.4 

38.4 

1292 

24.4 

24.3 

24.4 

1066 

'  59.5 

59.3 

61.8 

1362 

14.4 

16.4 

16.6 

1138 

55.0 

54.8 

57.0 

When  the  bands  are  as  wide  as  the  1/z  and  1.25/*  water-bands,  it 
will  be  seen  from  table  5  that  slit-widths  as  great  as  0.3  mm.  can  be 
employed  without  causing  any  appreciable  difference  in  the  intensity 
of  the  absorption.  The  slit  should,  however,  be  kept  as  narrow  as 
possible,  and  as  we  could  easily  make  fairly  accurate  measurements  in 
this  region  with  narrow  slits,  we  determined  to  use  in  all  of  the  work 
described  below  a  slit-width  of  0.2  mm. 


BY   MEANS   OF   THE   RADIOMICROMETER.  19 

THE  ABSORPTION  OF  LIGHT  BY  WATER  AS  AFFECTED  BY  HYDRATED  AND  BY 
NON-HYDRATED  SALTS. 

It  was  found  by  Guy,  Shaeffer,  and  Jones1  that  the  absorption  of 
light  by  water  was  changed  by  the  presence  of  hydrated  salts  which 
themselves  had  no  absorption.  Solutions  of  non-hydrated  salts  of 
equal  concentration  did  not  appreciably  affect  the  power  of  water  to 
absorb  light.  The  results  obtained  by  the  above-named  authors  were 
considered  by  them  as  strong  evidence  in  favor  of  the  solvate  theory  of 
solution  as  proposed  in  this  laboratory  about  15  years  ago. 

That  water  of  hydration  should  absorb  so  differently  from  an  equal 
amount  of  uncombined  water  was  regarded  as  a  fact  of  some  impor- 
tance, especially  in  its  bearing  on  the  solvate  theory  of  solution,  and 
one  deserving  very  careful  study.  Since  the  work  recorded  in  the  first 
paper  was  essentially  preliminary  in  character,  it  was  decided  to 
repeat  and  extend  this  work  with  our  improved  apparatus,  taking 
advantage  of  what  we  had  recently  learned  from  experience  in  con- 
nection with  such  work.  This  seemed  all  the  more  desirable  in  that 
our  earlier  results  had  been  interpreted  in  a  different  manner  by  Livens.2 
The  most  recent  work  gives  results  confirming  the  belief  that  our  origi- 
nal interpretation  was  correct. 

The  earlier  work  showed  clearly  that  concentrated  solutions  of  such 
hydrated  salts  as  calcium  chloride,  magnesium  chloride,  and  aluminium 
sulphate,  which  themselves  had  no  absorption  over  the  region  of  the 
1/z  and  the  1.25/z  water-bands,  had  a  very  different  transmission  curve 
from  that  of  a  layer  of  water  equal  in  depth  to  the  water  in  the  solution 
in  question.  For  some  wave-lengths  of  light  the  transparency  of  the 
solution  was  considerably  greater  than  that  of  an  equal  amount  of 
pure  water.  It  was  concluded  that  combined  water  has  kss  power  to 
absorb  light  than  free  water.  No  such  marked  effect  could  be  noted 
when  studying  the  absorption  of  aqueous  solutions  of  such  non- 
hydrated  salts  as  ammonium  chloride,  ammonium  nitrate,  and  potas- 
sium chloride,  even  when  the  solution  was  of  five  normal  concentra- 
tion. For  non-hydrated  salts  the  curves  for  the  solution  and  for  water 
were  nearly  identical.  It  was  therefore  clear  that  combined  water  had 
very  different  action  on  light  from  ordinary  free  water. 

We  have  repeated  practically  all  of  the  earlier  work  dealing  with 
this  phase  of  our  problem,  and  have  also  duplicated  a  very  large  part 
of  our  own  bearing  upon  it.  It  is  a  matter  of  some  difficulty  to 
avoid  considerable  errors  in  work  of  this  character,  and  for  this  reason 
it  was  considered  desirable  to  make  duplicate  measurements  as  often 
as  possible.  That  some  idea  may  be  gained  as  to  the  accuracy  with 
which  we  could  repeat  our  work,  we  give  in  a  number  of  cases  the 
results  of  such  duplicate  measurements. 

—  Carnegie  Inst.  Wash.  Pub.  No.  190.        Phys.  Zeit.,  14,  278  (1913).        *Ibid.,  14,  660  (1913). 


20  STUDY   OF   ABSORPTION   SPECTRA 

METHOD  OF  PROCEDURE. 

Suppose  we  wished  to  study  a  five  normal  solution  of  calcium  chloride. 
From  the  known  concentration  of  the  solution  and  its  specific  gravity, 
we  can  readily  calculate  the  volume  percentage  of  water  which  the 
solution  contains.  At  20°,  the  temperature  at  which  all  of  this  work 
was  done,  it  was  found  that  the  above-named  solution  contains  about 
90  per  cent  of  water.  The  1/x  water-band  is  of  such  intensity  that 
20  mm.  of  water  are  ample  for  its  study.  The  maximum  absorption  at 
the  center  of  the  band,  for  such  a  depth  of  layer,  is  about  60  per  cent. 
We  can  study  any  changes  better  when  the  absorption  is  not  so  intense 
as  to  give  small  deflections.  Furthermore,  the  deflections  given  by  the 
radiomicrometer  under  these  conditions  would  be  sufficient  to  allow 
accurate  readings.  In  all  of  this  work,  unless  otherwise  stated,  the 
width  of  the  slit  was  0.2  mm.,  and  the  current  0.8  ampere.  Under  these 
conditions  the  deflection  for  20  mm.  of  water  or  solution  would,  at  the 
point  of  maximum  absorption  or  the  center  of  the  band,  be  over  60  mm. 

The  absolute  transmission  for  20  mm.  of  the  solution  of  calcium 
chloride  would  be  determined  according  to  the  principle  already  out- 
lined. Its  transmission  curve  could  then  be  compared  directly  with 
that  for  18  mm.  of  pure  water.  A  layer  of  water  of  this  depth  absorbs 
so  intensely  at  1.25/i,  that  it  would  be  impossible  to  study  under  these 
conditions  the  intense  1.25ju  water-band.  If  10  mm.  of  solution  are 
used,  the  1.25/i  water-band  can  be  studied  very  satisfactorily,  the 
absorption  at  the  center  of  the  band  being  about  65  per  cent  and  the 
minimum  deflection  at  this  point  being  close  to  70  mm.  There  are, 
then,  four  transmission  curves  for  each  salt — one  for  the  solution  and 
one  for  the  solvent  as  the  result  of  studying  the  1/i  water-band,  and 
one  for  the  solution  and  one  for  the  solvent  when  studying  the  1.25// 
water-band. 

THE  RESULTS. 

Table  6  contains  the  results  of  two  series  of  measurements  with 
magnesium  chloride.  The  second  series  was  made  about  four  weeks 
after  the  first.  The  transmissions  are  those  of  20  mm.  of  the  solution. 

In  table  7  are  given  the  results  for  two  series  of  measurements  with 
potassium  chloride.  New  solutions  were  made  up  for  the  second  series 
of  measurements.  These  results  show  that  the  work  can  be  repeated 
with  a  fair  degree  of  concordance. 

Tables  6  and  7  also  show  that  there  is  a  marked  difference  between 
the  transmission  of  the  hydrated  salt  (magnesium  chloride)  and  a  layer 
of  water  equal  in  depth  to  that  in  the  solution.  No  such  difference 
exists  in  the  case  of  potassium  chloride,  a  typical  non-hydrated  salt. 

These  differences  come  out  most  clearly  for  those  wave-lengths  of 
light  which  are  most  strongly  absorbed.  Consider  the  region  about 
the  Iju  band,  table  6,  say  at  1.2/z,  where  the  deflections  are  fairly  large 


BY   MEANS   OF   THE   RADIOMICROMETER. 


21 


and  the  scale  can  be  accurately  read;  we  see  that  for  potassium  chloride 
the  difference  between  the  percentage  transmission  for  the  solution  and 
for  the  solvent  is  only  1  part  in  23,  whereas  for  magnesium  chloride, 
figure  1,  it  is  nearly  7  parts  in  26,  or  a  difference  of  about  25  per  cent. 
i&  Table  6  brings  out  another  point  of  considerable  interest.  If  we 
consider  the  center  of  the  l^t  band,  which  is  the  region  of  maximum 
absorption,  it  will  be  seen  that  for  potassium  chloride  the  solution 

TABLE  6. — Depth  of  cell  20  mm.    Magnesium  and  potassium  chlorides. 


X 

KC1 
3.5  N. 

H2O. 

MgCl2 
4.49  N. 

H20. 

MgCl2 
4.49  N. 

H20. 

706 

I  98.4 

98.0 

96.4 

97.3 

96.4 

96.4 

706 

I  98.4 

98.0 

95.3 

97.3 

97.3 

97.3 

746 

93.6 

96.0 

90.3 

94.9 

93.7 

93.7 

766 

92.5 

94.5 

93.7 

95.4 

92.3 

94.6 

770 

92.8 

94.8 

91.6 

95.1 

774 

92.0 

95.7 

92.8 

93.9 

778 

90.3 

94.3 

91.2 

93.1 

787 

93^5 

94!s 

92.0 

96.2 

94.4 

95.3 

809 

96.5 

96.5 

92.2 

95.8 

92.7 

95.1 

833 

93.3 

94.0 

93.0 

94.0 

93.6 

93.5 

857 

90.7 

92.1 

88.9 

91.8 

91.2 

91.6 

886 

90.0 

90.0 

88.5 

90.4 

90.4 

89.9 

913 

89.8 

88.5 

85.3 

88.7 

87.3 

88.3 

941 

81.5 

81.3 

80.6 

79.7 

82.3 

79.7 

952 

72.7 

72.3 

74.6 

70.0 

76.7 

71.4 

965 

63.7 

61.2 

65.6 

59.0 

67.4 

60.8 

972 

56.8 

56.3 

60.7 

54.4 

61.7 

54.7 

978 

51.8 

51.3 

56.2 

49.7 

56.6 

50.8 

984 

48.5 

47.8 

52.7 

46.6 

53.3 

47.3 

990 

45.2 

45.2 

49.4 

45.1 

49.6 

45.3 

996 

44.3 

45.0 

46.8 

44.7 

47.0 

44.6 

1003 

f  43.8 

44.6 

46.0 

44.9 

46.0 

45.5 

1003 

1  43.8 

45.0 

47.2 

45.3 

46.3 

45.3 

1010 

45.8 

45.8 

45.6 

46.8 

46.0 

46.5 

1017 

46.7 

47.5 

46.7 

48.3 

46.8 

48.4 

1024 

50.0 

50.0 

48.7 

51.3 

48.6 

50.3 

1029 

52.6 

52.5 

50.7 

53.7 

59.9 

52.9 

1035 

55.9 

55.2 

53.3 

57.0 

52.8 

55.3 

1041 

59.2 

58.7 

56.2 

60.0 

55.5 

58.4 

1053 

66.2 

64.8 

62.2 

66.3 

61.6 

64.3 

1066 

71.8 

70.3 

66.7 

71.5 

65.3 

68.7 

1081 

75.3 

73.2 

70.9 

74.3 

68.7 

73.3 

1095 

75.6 

74.8 

71.5 

75.7 

71.2 

74.3 

1110 

75.3 

74.3 

70.4 

74.2 

74.4 

73.4 

1123 

73.0 

71.5 

67.3 

70.2 

64.4 

71.2 

1138 

66.7 

65.8 

63.3 

64.1 

63.4 

64.8 

1146 

62.4 

61.0 

1153 

57.3 

56.0 

56.6 

53.7 

57.3 

55.2 

1161 

52.1 

49.6 

51.6 

47.4 

52.8 

50.0 

1168 

45.8 

44.7 

46.8 

41.7 

48.6 

43.0 

1175 

39.5 

38.3 

41.5 

34.8 

43.7 

37.0 

1183 

33.8 

32.5 

36.2 

29.5 

37.3 

31.8 

1192 

28.5 

27.2 

30.4 

24.5 

32.2 

26.4 

1200 

22.8 

21.9 

26.1 

19.6 

27.4 

21.6 

1208 

18.2 

17.1 

21.3 

16.1 

23.2 

17.4 

1216 

14.5 

14.7 

17.8 

14.2 

19.1 

14.5 

1224 

12.7 

12.7 

14.8 

12.6 

15.7 

12.6 

1231 

12.1 

12.1 

13.4 

11.8 

13.5 

11.8 

22 


STUDY   OF   ABSORPTION    SPECTRA 


absorbs  more  intensely  than  the  solvent.  Just  the  opposite  effect  is 
noted  for  magnesium  chloride,  in  which  case  the  solution  absorbs  less 
intensely.  The  same  relation  manifests  itself  again  in  the  case  of  the 
1.25/j  band,  as  can  be  seen  in  figure  2  and  in  table  7. 

The  curves  are  plotted  in  all  cases  from  the  data  given  in  the  tables. 
The  heavier  line  indicates  the  transmission  curve  for  the  solution;  the 
lighter  represents  that  for  the  solvent. 

Tables  6  and  7  are  plotted  as  curves  for  magnesium  chloride  in 
figures  1  and  2. 

TABLE  7. — Depth  of  cell  10  mm.     Potassium  and  magnesium  chlorides. 


X 

KC1 
3.5  N. 

H20. 

KC1 
3.5  N. 

H20 

MgCl2 
4.49  N. 

H20 

706 

/  97.0 

99.0 

98.3 

98.3 

98.3 

96.5 

706 

\  97.0 

99.0 

98.3 

98.3 

98.1 

96.5 

1153 

72.9 

73.6 

73.3 

73.5 

74.8 

72.7 

1168 

63.2 

63.4 

63.1 

63.4 

67.2 

62.3 

1175 

58.2 

58.3 

58.2 

58.3 

62.6 

57.4 

1183 

52.5 

53.6 

52.0 

53.5 

58.5 

52.9 

1192 

47.8 

48.5 

47.0 

48.5 

55.1 

48.7 

1200 

42.7 

44.4 

42.4 

44.5 

49.7 

44.3 

1208 

38.3 

40.2 

37.7 

40.0 

45.8 

40.0 

1216 

34.8 

37.2 

34.0 

37.4 

41.7 

37.0 

1231 

33.0 

34.5 

31.6 

34.6 

36.9 

34.0 

1238 

31.6 

33.2 

31.0 

34.2 

35.5 

33.8 

1246 

\  31.0 

33.2 

31.4 

34.1 

34.2 

33.3 

1246 

\  31.2 

33.5 

31.6 

34.0 

34.2 

33.3 

1254 

33.6 

33.8 

1261 

32.7 

33!3 

33.3 

34.2 

33.4 

34.3 

1269 

33.4 

34.4 

33.4 

35.4 

33.8 

34.6 

1278 

34.6 

34.6 

34.8 

34.3 

34.9 

35.4 

1292 

36.1 

35.1 

36.6 

35.4 

35.0 

35.7 

1306 

37.0 

35.1 

37.6 

35.2 

35.1 

35.4 

1313 

37.2 

34.7 

38.0 

35.1 

35.6 

35.2 

1321 

36.9 

33.4 

37.3 

33.8 

34.3 

34.3 

1334 

35.7 

32.3 

35.0 

32.3 

33.0 

32.2 

1348 

32.8 

29.1 

31.7 

28.9 

31.0 

28.8 

1362 

28.7 

24.4 

27.5 

24.2 

28.2 

24.5 

1377 

24.2 

19.7 

24.5 

20.2 

23.6 

20.4 

1388 

19.0 

15.8 

18.0 

15.8 

19.5 

15.4 

1400 

15.1 

12.5 

14.0 

12.2 

15.5 

11.9 

1414 

11.3 

10.0 

10.0 

9.5 

14.1 

11.7 

Tables  8  and  9  contain  the  percentage  transmissions  for  the  three 
hydrated  salts,  magnesium  bromide,  magnesium  sulphate,  and  zinc 
sulphate.  Where  the  light  is  intensely  absorbed  as  at  1.2/i,  and  the 
depth  of  layer  is  20  mm.,  the  solutions  are  over  30  per  cent  more  trans- 
parent than  the  solvent.  In  the  case  of  magnesium  chloride,  bromide, 
sulphate,  and  zinc  sulphate  at  the  centers  of  both  the  lju  and  the  1.25/x 
bands  (figs.  1  to  8)  the  solutions  are  the  more  transparent.  This 
difference  is  very  pronounced  in  the  case  of  the  two  sulphates,  especially 
with  magnesium  sulphate.  Here  the  difference  for  the  Iju  band  is 
about  18  per  cent,  and  for  the  1.25ju  band  about  20  per  cent. 


BY   MEANS   OF   THE   RADIOMICROMETER. 


23 


The  data  for  the  two  hydrated  salts  (zinc  nitrate  and  calcium 
chloride)  are  recorded  in  tables  10  and  11.  The  last  three  columns 
of  table  10  contain  triplicate  results  for  20  mm.  of  water,  while  the 
last  three  columns  in  table  11  contain  triplicate  results  for  10  mm.  of 
water.  To  save  time,  we  proposed  to  determine  carefully  the  trans- 
mission curve  for  20  mm.  of  water,  and  then  use  the  results  in  our 

TABLE  8. — Depth  of  cell  20  mm.    Transmission  of  hydrated  salts. 


X 

MgBr2 
3.60  N. 

H20. 

MgS04 
2.14  N. 

H20. 

ZnSO4 
2.92  N. 

H20. 

706 

/  98.0 

95.7 

96.5 

96.5 

96.4 

96.6 

706 

1  98.0 

95.7 

96.5 

96.7 

96.4 

96.6 

725 

.... 

94.4 

95.8 

95.3 

94.3 

746 

91.3 

94.'  8 

93.5 

92.7 

94.7 

94.0 

766 

95.5 

94.6 

93.5 

92.7 

94.2 

93.6 

774 

94.0 

778 

90.5 

92.8 

92.5 

94.2 

787 

.... 

92.5 

93!<5 

93.8 

95.4 

809 

9CK5 

93.6 

93.5 

96.0 

94.5 

95.2 

833 

91.5 

94.4 

93.0 

92.3 

92.3 

92.3 

857 

89.0 

91.0 

90.5 

89.8 

91.4 

91.0 

886 

86.6 

89.7 

89.0 

87.5 

89.2 

90.2 

913 

85.7 

87.3 

86.2 

85.3 

86.8 

87.4 

941 

78.8 

79.7 

78.6 

76.0 

79.6 

952 

72.8 

71.3 

72.3 

66.8 

74.2 

69.5 

965 

62.7 

58.6 

63.4 

54.8 

64.8 

58.1 

972 

58.4 

48.8 

59.2 

52.7 

978 

52^6 

48  '.2 

53.3 

44.3 

65.2 

47.7 

984 

47.6 

45.3 

50.0 

40.3 

51.0 

44.0 

990 

45.7 

43.5 

47.2 

38.6 

48.6 

42.4 

996 

'  44.7 

42.8 

45.3 

37.1 

46.4 

41.4 

996 

45.3 

36.9 

45.8 

41.3 

1003 

f  45.2 

43^3 

44.6 

36.9 

45.6 

41.6 

1003 

[44.3 

43.3 

1010 

45.6 

44.8 

45.2 

38.3 

45.9 

43.6 

1017 

46.0 

46.4 

46.2 

40.0 

46.6 

45.0 

1024 

48.3 

48.8 

48.4 

42.7 

48.7 

47.3 

1029 

51.7 

51.6 

49.4 

45.2 

50.7 

50.4 

1035 

55.0 

54.6 

52.0 

52.9 

1041 

57.0 

58.0 

55.5 

51.8 

54.7 

56.5 

1053 

66.6 

64.6 

59.8 

58.7 

59.2 

64.3 

1066 

70.0 

69.3 

65.2 

63.6 

62.8 

68.1 

1081 

72.3 

73.0 

67.8 

68.0 

66.3 

71.4 

1095 

75.8 

74.6 

70.0 

70.3 

68.4 

72.3 

1110 

73.5 

73.6 

69.3 

68.4 

67.6 

71.0 

1123 

73.0 

69.2 

66.5 

65.5 

1138 

67.6 

63.9 

62.0 

59.8 

62.5 

6l!2 

1146 

59.6 

56.4 

1153 

61.0 

53.5 

54.5 

49!4 

55.8 

51.2 

1161 

.... 

51.7 

45.4 

1168 

53  .B 

39.4 

44^3 

38.8 

46.4 

39.1 

1175 

45.4 

32.6 

38.7 

32.3 

40.8 

32.7 

1183 

.... 

33.2 

26.2 

35.6 

27.1 

1192 

siie 

21.1 

28.3 

21.3 

30.3 

21.8 

1200 

25.2 

17.0 

23.1 

16.3 

25.3 

19.3 

1208 

19.6 

14.1 

19.4 

12.9 

21.3 

13.9 

1216 

16.4 

12.0 

15.4 

10.1 

17.7 

11.4 

1224 

15.2 

11.4 

13.3 

8.4 

15.1 

10.2 

1231 

12.8 

10.6 

11.5 

7.3 

13.0 

9.3 

24 


STUDY   OF   ABSORPTION    SPECTRA 


subsequent  work.  The  cells  containing  the  solution  could  then  be  so 
adjusted  that  they  would  contain  just  as  much  water  as  was  present 
when  the  percentage  transmissions  for  20  mm.  of  water,  that  is  21  mm. 
in  one  cell  and  1  mm.  in  the  other,  were  determined. 

It  was  found  that,  although  this  procedure  would  save  much  tune 
and  labor,  it  can  not  be  recommended  where  it  is  desired  to  do  accurate 
work,  since  the  conditions  would  not  be  as  nearly  comparable  as  they 
would  be  if  the  absorption  of  the  solvent  was  determined  immediately 

TABLE  9. — Depth  of  cell  10  mm.    Transmission  of  hydrated  satis. 


X 

MgBr2 
3.60  N. 

H20. 

MgSO< 
2.14  N. 

H20. 

ZnSO4 
2.92  N. 

H20. 

706 

f  97.0 

98.4 

96.5 

96.6 

99.0 

97.0 

706 

\  97.3 

98.4 

96.5 

96.6 

99.0 

97.0 

1138 

79.1 

77.3 

78.6 

76.1 

1153 

73.6 

69.7 

72.6 

69.0 

M.'.t 

75.Q 

1168 

66.2 

59.8 

64.8 

58.0 

74.6 

65.6 

1175 

61.2 

54.1 

59.0 

52.3 

70.3 

60.0 

1183 

55.8 

48.3 

55.5 

47.1 

65.3 

55.3 

1192 

51.5 

44.1 

49.6 

42.2 

59.7 

51.3 

1200 

46.4 

39.3 

45.4 

36.8 

56.4 

46.6 

1208 

41.6 

36.6 

41.4 

32.8 

50.5 

43.6 

1216 

38.9 

33.7 

38.2 

29.8 

47.5 

40.4 

1224 

36.2 

33.0 

35.6 

27.8 

1231 

34.4 

31.8 

34.6 

27.1 

42^6 

39!  i 

1238 

33.3 

31.3 

32.9 

26.0 

41.7 

38.1 

1246 

f  33.3 

31.2 

32.1 

25.9 

41.4 

38.2 

1246 

\  32.8 

31.1 

32.4 

25.7 

41.7 

38.1 

1254 

31.3 

31.6 

31.7 

25.9 

41.7 

38.5 

1261 

34.3 

31.7 

31.5 

26.0 

42.2 

39.3 

1269 

34.4 

32.0 

31.3 

26.4 

42.7 

39.7 

1278 

34.8 

32.2 

30.9 

26.6 

43.3 

40.3 

1292 

36.2 

32.3 

30.8 

27.5 

43.7 

40.7 

1306 

37.6 

32.8 

30.3 

27.9 

43.7 

39.7 

1313 

37.0 

32.4 

30.1 

26.9 

43.1 

39.1 

1321 

37.7 

31.3 

29.2 

26.7 

42.7 

36.4 

1334 

35.8 

28.8 

27.5 

25.1 

37.8 

32.6 

1348 

32.5 

25.1 

25.2 

22.8 

33.4 

27.4 

1362 

27.5 

20.6 

19.9 

19.2 

28.7 

23.3 

1377 

22.8 

16.4 

17.4 

15.7 

22.9 

17.5 

1388 

17.7 

12.2 

13.8 

11.7 

17.4 

13.1 

1400 

13.4 

9.0 

10.1 

9.4 

13.6 

9.7 

1414 

9.0 

6.0 

7.9 

6.0 

9.6 

6.9 

1429 

6.3 

5.3 

5.2 

3.3 

after  that  of  the  solution.  The  latter  procedure  was  therefore  adopted. 
It  will  be  noted  that  the  percentage  transmissions  for  the  three  series 
of  measurements  with  pure  water  (tables  10  and  11)  are  very  concord- 
ant for  both  the  1/z  and  the  1.25ju  bands. 

Zinc  nitrate,  unlike  the  hydrated  chlorides,  bromides,  and  sulphates, 
presents  an  apparent  exception  in  the  case  of  the  IJJL  band  (fig.  9).  At 
the  center  of  this  band  the  solution  has  the  greater  absorption,  although 
the  difference  between  the  absorption  of  the  solution  and  that  of  the 
solvent  is  not  large.  The  same  phenomenon  makes  its  appearance  in 


BY   MEANS   OF   THE    RADIOMICROMETER. 


25 


the  case  of  magnesium  nitrate,  the  work  having  been  three  times 
repeated  and  the  same  results  obtained  in  every  case.  We  desire  to 
arrest  attention  to  this  apparent  peculiarity  of  the  hydrated  nitrates. 
The  1.25ju  band  (fig.  10)  does  not  show  this  peculiarity.  The  results 
here  are  of  the  same  general  character  as  in  the  cases  of  the  other 
hydrated  salts  with  which  we  worked.  There  are  differences  as  great  as 
25  per  cent  between  the  absorption  of  the  solution  and  that  of  the  sol- 


-- 


MgCI2,4.49N 
Depth  20mm. 


0.9  0.95 


1.0  1.05  1.10 

FIG.  1. — Wave-lengths. 


H20 


M§CI2,449 
Depth  10mm 


1.1 


1.20  1.25  1.30 

FIG.  2. — Wave-lengths. 


26 


STUDY   OF   ABSORPTION   SPECTRA 


vent  in  the  neighborhood  of  1.2/j,  (fig.  9)  where  the  absorption  is  intense, 
as  it  is  approaching  the  center  of  the  1.25/j  water-band.  Here  the  solu- 
tion is  the  more  transparent,  as  is  generally  the  case  with  hydrated  salts. 
The  curves  for  calcium  chloride  (figs.  11  and  12)  deserve  careful 
study.  The  solution  is  more  transparent  than  the  solvent  for  every 
wave-length  of  light  that  we  studied.  This  holds  true  for  both  of  the 

TABLE  10. — Depth  of  cell  20  mm.     Transmissions  of  zinc  nitrate, 
calcium  chloride,  and  water. 


X 

Zn(NO,)2 
4.30  N. 

H20. 

CaCl2 

4.78  N. 

H20. 

H20. 

H20. 

H20. 

706 

f  99.0 

94.1 

97.1 

96.3 

96.5 

96.5 

96.3 

706 

\  99.0 

95.6 

97.1 

96.4 

96.5 

96.5 

96.0 

746 

99.0 

90.3 

96.9 

96.2 

766 

95.6 

90.9 

95.6 

95.5 

770 

95.8 

92.9 

94.1 

96.7 

774 

96.0 

93.0 

94.1 

94.6 

778 

97.2 

94.8 

95.2 

95.7 

787 

96.2 

94.4 

93.8 

96.1 

809 

96.7 

94.4 

95.3 

95.7 

833 

96.2 

93.6 

95.3 

94.6 

857 

94.6 

93.7 

91.9 

91.9 

886 

92.6 

88.8 

91.3 

90.9 

9l!i 

91  A 

90.7 

913 

88.7 

88.0 

89.7 

88.1 

87.9 

88.1 

87.9 

941 

82.7 

79.2 

82^7 

79.3 

78.5 

78.7 

79.5 

952 

77.5 

72.4 

75.7 

70.9 

69.9 

70.3 

70.8 

965 

67.8 

61.0 

66.7 

59.5 

58.2 

58.4 

59.5 

972 

62.6 

57.3 

60.7 

54.5 

53.0 

54.0 

54.8 

978 

57.3 

54.3 

54.9 

50.3 

49.6 

49.3 

49.7 

984 

52.3 

50.7 

50.7 

46.6 

46.8 

47.0 

46.4 

990 

48.7 

48.2 

47.3 

44.4 

44.7 

46.3 

44.7 

996 

47.2 

47.3 

44.8 

43.5 

44.4 

43.9 

43.5 

1003 

f  45.6 

47.6 

44.3 

43.7 

45.2 

45.1 

44.5 

1003 

1  45.7 

47.3 

44.3 

42.4 

44.8 

44.9 

44.4 

1010 

46.2 

48.7 

44.7 

44.4 

46.2 

46.1 

46.3 

1017 

48.3 

50.6 

46.8 

45.8 

47.8 

47.8 

47.7 

1024 

50.4 

51.4 

48.7 

48.3 

50.2 

49.7 

49.4 

1029 

53.2 

55.5 

52.0 

51.2 

53.7 

52.7 

52.6 

1035 

55.8 

57.7 

55.4 

54.1 

56.2 

56.3 

55.2 

1041 

58.8 

60.7 

61.7 

57.4 

59.3 

58.9 

60.3 

1053 

68.8 

66.8 

66.5 

63.2 

65.3 

66.2 

65.2 

1066 

68.8 

70.6 

71.8 

68.9 

70.7 

70.6 

69.9 

1081 

72.2 

73.7 

74.7 

71.6 

74.6 

73.3 

73.5 

1095 

73.3 

75.1 

76.2 

71.3 

76.2 

75.0 

75.0 

1110 

74.0 

73.7 

74.5 

71.2 

74.5 

73.3 

72.7 

1123 

74.3 

70.8 

72.4 

67.2 

71.3 

70.3 

70.7 

1138 

69.3 

65.7 

67.7 

61.0 

64.1 

63.5 

64.0 

1146 

.... 

64.7 

55.5 

59.5 

58.3 

58.9 

1153 

64!s 

56.3 

59.8 

49.7 

53.6 

52.8 

53.3 

1161 

59.0 

50.3 

54.9 

43.3 

47.7 

47.2 

47.1 

1168 

53.8 

44.7 

48.7 

37.3 

41.7 

40.7 

43.6 

1175 

47.7 

38.0 

42.5 

30.4 

35.2 

35.2 

34.3 

1183 

42.0 

32.9 

37.4 

24.4 

30.5 

29.2 

29.7 

1192 

36.7 

28.1 

30.9 

19.9 

25.2 

24.3 

24.2 

1200 

30.4 

23.2 

25.7 

15.0 

20.0 

19.8 

19.5 

1208 

25.2 

19.2 

21.3 

13.1 

16.7 

15.7 

16.0 

1216 

21.1 

16.4 

17.0 

11.3 

13.9 

13.3 

13.5 

1224 

18.4 

15.1 

14.3 

10.9 

1231 

16.2 

14.7 

12.9 

10.2 

BY   MEANS   OF   THE   RADIOMICROMETER. 


27 


water-bands,  and  there  is  no  evidence  which  indicates  that  a  shift  has 
taken  place,  as  is  apparently  the  case  with  a  few  of  the  other  bands  for 
hydrated  salts.  Whenever  there  is  a  shift  of  the  band  for  a  hydrated 
salt,  it  is  always  towards  the  red.  Near  the  end  of  the  1/z  water-band 
for  calcium  chloride  (fig.  11),  the  solution  is  about  40  per  cent  more 
transparent  than  the  solvent.  This  is  the  greatest  difference  found 
with  any  of  the  salts  that  we  studied.  Near  1 .3/i  (fig.  12) ,  which  is  close 
to  the  1.25ju  water-band,  the  solution  is  over  35  per  cent  more  trans- 


X 


MgBr2,3.60N 
Depth  20mm. 


0.95 


1.05  1.10 

FIG.  3.— Wave-lengths. 


MgBr2,3.60N 
Deplh  Ipmm. 


1.20  1.25  1.30 

Fro.  4. — Wave-lengths. 


1.35 


28 


STUDY   OF   ABSORPTION    SPECTRA 


parent  than  the  solvent.  No  such  differences  were  found  at  or  near  the 
1.25/*  water-band  for  any  of  the  other  hydrated  salts  that  were  studied. 
The  solution  of  the  calcium  chloride  used  was  nearly  5  times  normal,  and 
it  was  found  that  the  above-named  differences  increase  with  increase  hi 
concentration.  This  was  expected,  since  the  total  amount  of  combined 
water  increases  with  the  concentration  of  the  solution. 

The  first  column  of  table  12  contains  the  percentage  transmissions 
for  20  mm.  of  water,  while  that  of  table  13  gives  the  transmissions  for 
10  mm.  of  water.  The  figures  given  are  the  averages  of  several  series 

TABLE  11. — Depth  of  cell  10  mm.     Transmission  of  zinc  nitrate,  calcium 
chloride,  and  water. 


\ 

Zn(N03)2 
4.30  N. 

H20. 

CaClj 

4.78  N. 

H20. 

H20. 

H20. 

H20. 

706 

f  99.0 

97.0 

98.6 

98.0 

98.3 

98.5 

98.3 

706 

\  99.0 

97.0 

98.6 

98.0 

98.3 

98.5 

98.3 

1153 

84.9 

75.6 

75.6 

71.2 

71.3 

70.5 

70.7 

1168 

76.4 

65.6 

66.8 

61.3 

61.0 

59.8 

59.7 

1175 

70.3 

60.3 

61.3 

55.8 

55.3 

54.2 

54.0 

1183 

65.3 

55.3 

56.2 

50.3 

50.0 

48.6 

49.0 

1192 

59.7 

51.3 

52.6 

45.5 

45.3 

44.2 

44.5 

1200 

56.4 

46.6 

46.7 

41.3 

40.7 

39.8 

40.0 

1208 

50.5 

43.6 

42.7 

37.6 

37.2 

36.9 

37.2 

1216 

47.5 

40.4 

39.4 

34.9 

35.4 

35.1 

34.9 

1231 

42.6 

39.1 

34.3 

33.6 

33.8 

33.4 

33.1 

1238 

41.7 

38.1 

34.1 

32.8 

33.2 

33.0 

33.2 

1246 

f  41.4 

38.2 

33.9 

32.4 

33.7 

33.4 

33.5 

1246 

I  41.7 

38.2 

33.9 

32.4 

32.6 

33.7 

33.4 

1254 

41.7 

38.5 

34.3 

32.4 

34.0 

34.2 

33.6 

1261 

42.2 

38.8 

34.9 

32.9 

34.3 

34.1 

33.8 

1269 

42.7 

39.3 

36.1 

33.3 

1278 

43.3 

39.7 

37.1 

34.0 

35.7 

34.9 

33.7 

1292 

43.3 

40.3 

39.4 

34.7 

35.8 

35.7 

35.3 

1306 

43.7 

40.7 

40.5 

34.5 

35.5 

35.1 

35.  1 

1313 

43.7 

39.7 

1321 

43.1 

39.1 

ii.3 

32!4 

3.L5 

32^8 

34!2 

1334 

42.7 

36.4 

39.1 

29.8 

30.4 

30.4 

31.0 

1348 

37.8 

32.6 

36.1 

26.0 

25.7 

25.8 

25.7 

1362 

33.4 

27.4 

31.6 

20.6 

21.0 

21.1 

21.8 

1377 

28.7 

23.3 

26.2 

16.3 

16.2 

16.3 

17.2 

1388 

22.9 

17.5 

20.9 

12.2 

11.6 

12.8 

12.3 

1400 

17.4 

13.1 

15.1 

8.9 

1414 

13.6 

9.7 

12.1 

6.3 

of  measurements.  Three  of  these  series  are  given  in  tables  10  and  11 
These  transmissions  of  the  solvent  are  to  be  compared  with  the  two 
series  of  results  for  magnesium  nitrate  and  sodium  nitrate,  and  with 
that  for  ammonium  bromide.  A  sufficient  depth  of  solution  was  used 
to  contain  21  mm.  of  water  in  one  cell  and  1  mm.  of  water  in  the  second 
cell,  these  being  the  depths  of  layers  employed  in  determining  the 
percentage  transmissions  for  20  mm.  of  the  solvent.  In  obtaining  the 
results  recorded  in  table  13,  one  cell  had  enough  solution  to  contain 
11  mm.  of  water,  while  the  second  had  enough  solution  to  contain 


BY   MEANS   OF   THE    RADIOMICROMETER. 


29 


1  mm.  of  water.     This  procedure  saved  in  this  case  the  necessity  of 
making  a  series  of  measurements  with  the  solvent. 

Magnesium  nitrate  presented  so  many  peculiarities  that  the  work 
with  this  salt  was  repeated  three  times,  starting  in  every  case  with  a 
new  solution.  The  same  general  results  were  always  obtained.  The 
curve  for  the  I/*  band  (fig.  13)  was  plotted  from  data  not  recorded  in 
this  paper.  The  concentration  of  the  solution  was  3.80  normal,  and, 
as  in  other  cases,  21  mm.  of  solution  were  placed  in  one  cell  and  1  mm. 


\ 


Mg  S04,2.l4N 
Depth  20  mm. 


1.0  1.05  1.10 

FIG.  5. — Wave-lengths. 


1.20  1.25  1.30 

FIG.  6.— Wave-lengths. 


30  STUDY   OF   ABSORPTION    SPECTRA 

in  the  other.  The  solution  curve  for  the  1.25/i  band  (fig.  14)  was 
plotted  from  the  data  given  in  the  second  column  of  table  13. 

Magnesium  nitrate  (like  zinc  nitrate,  the  other  hydrated  nitrate 
which  we  studied)  shows  the  peculiarity  that  at  the  center  of  the  In 
band  (fig.  13)  the  solution  is  less  transparent  than  the  solvent.  The 
difference  in  this  case  is  quite  pronounced.  At  the  center  of  the  1.25;* 
band  (fig.  14)  the  solution  and  solvent  have  about  the  same  trans- 
mission, whereas  the  absorption  of  zinc  nitrate  (fig.  10)  at  this  point 
conforms  to  the  general  characteristics  of  the  bands  for  the  other 
hydrated  salts.  Magnesium  nitrate  presents  another  curious  feature, 
which  is  not  present  in  the  results  for  any  of  the  other  hydrated  salts. 
The  difference  between  the  percentage  transmissions  for  the  solution 
and  for  the  solvent  is  at  no  point  very  great.  The  data  given  in  table 
12  for  the  1/x  band  show  that  for  wave-lengths  of  light  near  1.2/x  the 
solution  is  actually  several  per  cent  less  transparent  than  the  solvent. 
This  is  the  region  where  solutions  of  hydrated  salts  in  general  are  much 
more  transparent  than  the  pure  solvent.  This  is  clearly  shown  in  both 
series  of  measurements  given  in  table  12.  Here  we  have  employed 
a  deeper  layer  of  solution  than  in  the  series  from  which  the  curve  for 
the  Ijj,  band  was  plotted.  Consequently,  the  intensity  of  absorption 
is  greater,  and  this  would  have  the  effect  of  making  more  pronounced 
any  peculiarity  shown  by  a  more  shallow  layer.  The  curves  for  the 
solution  and  solvent  cross  at  a  point  closer  to  the  violet  end  of  the 
spectrum  when  deeper  layers  are  used.  This  is  evident  from  data 
given  in  table  12,  and  the  results  clearly  show  that  the  solution  would 
not  have  the  greater  transparency  at  the  center  of  the  1.25ju  band. 

The  1.25ju  band  (fig.  14)  does  not  show  any  special  peculiarities, 
there  being  several  other  bands  of  very  similar  character  in  the  cases 
of  the  other  hydrated  salts.  As  might  be  expected  from  the  nature  of 
the  results  for  the  1/z  band,  the  solution  is  less  transparent  just  before 
the  center  of  the  band  is  reached. 

The  transmission  curves  for  ammonium  bromide  (figs.  15  and  16), 
being  so  widely  different  from  those  of  the  solvent,  have  been  plotted 
to  bring  out  the  differences  in  general  between  hydrated  and  non- 
hydrated  salts.  Ammonium  bromide  is  a  typical  non-hydrated  salt, 
and  the  results  for  this  salt  can  be  compared  with  those  for  magnesium 
bromide,  a  typical  hydrated  salt. 

In  every  case  thus  far  studied,  the  solution  of  a  non-hydrated  salt 
is  less  transparent  at  the  center  of  the  band  than  the  solvent.  The 
differences  in  the  case  of  ammonium  bromide  are  quite  marked.  It  is 
only  through  a  narrow  range  of  wave-lengths  (that  is,  from  1.1 6ju  to 
1.1 8/x,  fig.  13)  that  the  solution  is  the  more  transparent  in  the  neighbor- 
hood of  the  Ifj.  band,  and  even  here  the  difference  is  very  slight. 

Near  1.2/*,  where  the  hydrated  salts  are  much  more  transparent  than 
the  solvent,  ammonium  bromide  is  about  30  per  cent  less  transparent. 


BY   MEANS   OF   THE   RADIOMICROMETER. 


31 


The  curve  for  the  1.25/z  band  (fig.  16)  brings  out  exactly  the  same 
relations  as  have  been  noted  for  the  1/z  band.  These  curves  cross  in 
two  places,  as  would  be  expected;  first,  in  the  region  already  men- 
tioned for  the  I/*  band;  second,  for  a  small  region  around  1.3/u.  For 
all  other  wave-lengths  the  solution  is  less  transparent  than  the  solvent. 
Sodium  nitrate  gave  results  of  a  slightly  different  character  from 
any  that  had  been  obtained  for  non-hydrated  salts.  The  measure- 
ments were  repeated  three  times,  two  series  of  results  being  recorded. 
In  every  case  new  solutions  were  used,  yet  results  of  the  same  general 


ZnS04,2.92N 
Depth  20mm 


1.05  1.10 

FIG.  7.— Wave-lengths. 


K 

\ 

\ 

s 

\ 

!S 

n  -•    -= 

^^ 

\ 

ZnS04,2.92N 
Depth  10mm. 

\ 

^^ 

1.1                       1.15                         1.20                        1.25                         1.30                        1.35                         1.40 
FIG.  8.  —  Wave-lengths. 

32 


STUDY   OF   ABSORPTION    SPECTRA 


character  were  obtained  in  each  series  of  measurements.  The  curves 
for  sodium  nitrate  (figs.  17  and  18)  were  plotted  from  the  results  given 
in  column  5,  tables  12  and  13. 

The  solution  of  sodium  nitrate,  as  is  the  case  with  all  the  other 
non-hydrated  salts,  is  less  transparent  than  the  solvent  at  the  center 
of  both  bands.  The  difference  is  quite  marked  for  this  salt.  The 
solution  curve  for  both  bands  is  very  much  more  elongated  than  that 
for  the  solvent.  At  the  point  of  maximum  absorption  it  shows  the 
greater  absorption,  and  at  the  point  of  maximum  transparency  the 
greater  transparency. 

TABLE  12. — Magnesium  nitrate,  ammonium  bromide  and 
sodium  nitrate. 


A 

H20 

Mg(N03)2 
3.68  N. 

(MgN03)2 
3.63  N. 

NH<Br 

4.016  N 

NaNO3 
5.69  N. 

NaNOs 
4.95  N. 

709 

f  96.3 

96.5 

96.4 

95.6 

96.6 

96.3 

709 

\  96.3 

96.5 

96.4 

95.8 

96.6 

96.3 

890 

90.8 

89.8 

90.0 

88.0 

93.6 

87.8 

918 

88.0     87.3 

86.8 

84.8 

88.3 

88.2 

934 

85.0 

83.7 

84.0 

81.5 

84.6 

85.6 

946 

79.2 

80.0 

79.6 

76.8 

78.7 

79.8 

952 

75.4 

76.7 

76.0 

73.2 

74.7 

76.3 

957 

70.2  I   72.2 

71.2 

69.6 

68.7 

71.5 

963 

64.5  :   67.3 

66.4 

58.7 

63.0 

64.5 

977 

53.5 

54.3 

53.5 

50.7 

50  4 

51.0 

983 

49.5 

47.8 

47.3 

44.8 

45.5 

44.1 

989 

46.8 

42.7 

42.7 

40.3 

41.7    39.7 

995 

44.7 

39.6 

39.3 

38.6 

39.0 

37.3 

1002 

44.0 

38.5 

38.5 

37.0 

38.9 

36.4 

1008 

\  44.8 

38  7 

39.1 

37.0 

40.0 

37.5 

1008 

I  44.7 

39.2 

38.6 

37.7 

39.9 

36.2 

1014 

46.2 

40.8 

41.0 

38.4 

41.7 

39.6 

1022 

47.8 

43.6 

43.3 

40.3 

45  0 

42.0 

1027 

49.7 

46.3 

47.0 

43.3 

49.2 

46.4 

1035 

52.7 

50.0 

50.6 

46.1 

52.8 

51.2 

1042 

56.2 

54.2 

54.2 

49.2 

56.8 

55.2 

1048 

59.5 

57.7 

58.5 

52.6 

60.9 

59.4 

1060 

65.3 

64.5 

64.5 

58.2 

68.2 

67.3 

1073 

70.6 

68.6 

69.6 

62.8 

73.9 

73.8 

1088 

73.5 

72.8 

72.7 

63.6 

77.2 

77.7 

1103 

75.2 

74.7 

73.4 

65.5 

77.3 

78.2 

1118 

73.3 

73.3 

71.8 

65.6 

75.2 

78.2 

1132 

70.7 

70.2 

69.5 

63.3 

71.6 

73.7 

1139 

67.5 

68.9 

67.0 

61.4 

68.3 

71.3 

1147 

64.0 

65.7 

64.3 

60.3 

64.3 

67.7 

1154 

58.9 

62.2 

62.0 

57.2 

59.2 

64.4 

1162 

53.0 

58.2 

56.4 

53.5 

53.4 

58.4 

1170 

47.2 

52.7 

52.7 

48.1 

46.7 

52.7 

1177 

41.2 

44.7 

44.5 

41.2 

40.3 

44.5 

1186 

35.2 

36.9 

35.2 

33.1 

33.5 

34.8 

1195 

29.7 

28.8 

28.7 

25.4 

28.2 

26.4 

1203 

24.3 

21.9 

21.0 

18.4 

22.2 

18.5 

1210 

19.8 

15.9 

16.5 

13.5 

17.4 

12.8 

1220 

16.0 

12.6 

12.5 

10.0 

13.9 

9.6 

1227 

13.5 

11.0 

10.1 

11.2 

8.5 

BY   MEANS   OF   THE   RADIOMICROMETER. 
DISCUSSION  OF  THE  RESULTS. 


33 


The  chief  points  of  interest  brought  out  by  a  study  of  the  transmis- 
sion curves  for  hydrated  and  for  non-hydrated  salts  are  the  following : 
Aqueous  solutions  of  hydrated  salts  generally  have  greater  transpar- 
ency than  pure  water  at  the  centers  of  the  absorption  bands.  The 
exceptions  are  the  lyu  band  for  zinc  nitrate  (fig.  9)  and  magnesium 


H,0 


Zn(N03)2,4.30N 
Depth  20mm. 


1.0  1.05  1.10 

FIG.  9. — Wave-lengths. 


1.15 


1.20 


\ 


\ 


Zn(N03)  .4.301 
Depth  10mm. 


1.25  1.30 

FIG.  10. — Wave-lengths. 


34 


STUDY   OF   ABSORPTION    SPECTRA 


nitrate  (fig.  13)  and  the  1.25/j  band  for  magnesium  nitrate  (fig.  14). 
The  above  relation  is  shown  very  clearly  in  many  cases  by  a  study  of 
the  curves  for  hydrated  salts.  With  increasing  wave-lengths,  as  the 
regions  of  intense  absorption  are  approached  the  solution  is  much  more 
transparent  than  the  solvent.  Near  1.2/z,  as  we  approach  the  1.25ju 
band,  having  in  the  path  of  the  light  20  mm.  of  the  solution,  the  solu- 
tion is  usually  from  20  to  40  per  cent  more  transparent  than  the  pure 
solvent.  This  difference  increases  as  the  depth  of  the  solution  in  the 
path  of  the  light  increases.  It  is  very  much  less  for  10  mm.  of  solution 

TABLE  13. — Magnesium  nitrate,  ammonium  bromide 
and  sodium  nitrate. 


X 

H,0. 

Mg(N03)j 
3.68  N. 

Mg(NO,)2 
3.63  N. 

NH4Br  NaNO3 
4.016  N.j  5.69  N. 

NaN03 
4.95  N. 

709 

f  98.5 

98.3 

98.0 

98.0 

98.5 

98.0 

709 

[  98.5 

98.3 

98.0 

98.0  |  98.5 

98.3 

1147 

78.5 

82.8 

82.7 

78.3  I  79.0 

82.5 

1154 

75.2 

80.4 

80.3 

76.2 

75.5 

80.0 

1162 

70.6 

77.1 

77.0 

73.3 

70.7 

76.5 

1170 

65.0 

72.9 

72.7 

69.2 

64.8 

72.3 

1177 

59.7 

67.0 

66.4 

63.4 

59.7 

64.0 

1186 

54.5 

59.4 

59.4 

55.7 

52.7 

56.2 

1195 

49.0 

52.2 

52.0 

48.6 

47.7 

47.5 

1203 

44.7 

44.7 

44.7 

41.6 

42.3 

40.1 

1210 

40.0 

39.7 

39.0 

35.6 

37.5 

33.9 

1220 

37.2 

35.8 

34.5 

31.8 

34.2 

30.6 

1227 

33.8 

33.6 

33.0 

29.4 

32.0 

28.9 

1235 

33.4 

33.4 

32.7 

28.2 

30.8 

29.0 

1242 

33.3 

33.5 

33.0 

27.4 

30.4 

29.2 

1250 

33.1 

33.9 

33.1 

27.7 

31.6 

30.4 

1258 

f  33.2 

34.7 

33.6 

27.8 

31.2 

30.4 

1258 

\  33.4 

34.6 

33.6 

28.1 

31.2 

30.6 

1266 

33.6 

35.6 

34.3 

28.2 

33.6 

32.1 

1272 

34.0 

36.3 

35.0 

28.8  i  34.6 

33.6 

1289 

34.8 

37.6 

36.9 

30.0 

36.8 

36.8 

1303 

35.4 

39.3 

38.2 

31.1 

38.6 

38.3 

1318 

35.0 

39.3 

38.6 

31.9    38.7 

40.0 

1331 

33.4 

38.2 

38.4 

31.7  !  37.2 

39.5 

1345 

30.7 

36.0 

36.2 

30.0  i  34.7 

37.4 

1359 

25.8 

31.2 

31.5 

27.0  1  29.7 

31.9 

1373 

21.8 

25.0 

26.3 

21.8    24.6 

25.1 

1386 

16.7 

19.0 

21.0 

16.5    19.2 

19.5 

1399 

15.3 

12.7 

14.1 

11.5    15.2 

12.6 

than  for  20mm.,  for  the  same  wave-length  of  light.  It  requires  an  intense 
absorption  to  bring  out  these  differences  in  a  pronounced  manner.  A 
decrease  in  the  concentration  of  the  solution  also  decreases  the  differ- 
ence between  the  absorption  for  the  solution  and  for  the  solvent,  as  has 
been  pointed  out  by  Jones  and  Guy.1  || 

Non-hydrated  salts,  under  similar  conditions,  give  results  in  many 
respects  exactly  the  opposite  of  those  obtained  with  hydrated  salts. 
Three  non-hydrated  salts  were  investigated,  and  in  all  three  cases  the 


Carnegie  Inst.  Wash.  Pub.  No.  190,  page  49. 


BY   MEANS    OF   THE    RADIOMICROMETER. 


35 


solution  had  greater  absorption  than  the  solvent  at  the  centers  of  the 
bands.  In  the  previous  work  in  this  field,  by  Jones  and  Guy,1  two  other 
non-hydrated  salts  were  studied  (ammonium  chloride  and  ammonium 
nitrate) ,  and  these  showed  exactly  the  same  phenomenon.  It  will  be 
remembered  that  for  hydrated  salts  the  solution  was  more  transparent 
in  all  cases  except  the  nitrates.  The  transmissions  for  the  solution  of 
potassium  chloride  are  very  nearly  identical  with  those  of  water  in  the 


\\ 


\ 


CaCl2,4.78N 

Depth  20  mm 


1.0  1-05  1.10 

FIG.  11. — Wave-lengths. 


1.20 


CaCl2,478N 

Depth  lOmm. 


1.25  1.30 

Fio.  12.— Wave-lengths. 


1.40 


'Carnegie  Inst.  Wash.  Pub.  No.  190;  Phys.  Zeit.,  14,  279  (1913). 


36  STUDY   OF   ABSORPTION   SPECTRA 

case  of  the  I/*  band,  table  6;  therefore,  no  curve  was  plotted  for  this 
salt.  The  results  for  the  1.25/z  band,  table  7,  show  differences  between 
the  absorption  of  the  solution  and  of  the  solvent.  The  solution  has  the 
greater  absorption  at  the  center  of  this  band,  which  is  in  accord  with 
the  results  in  general  for  non-hydrated  salts.  Otherwise  the  results  are 
of  the  same  general  character  as  those  for  a  very  weakly  hydrated  salt. 

The  curves  for  ammonium  bromide  and  sodium  nitrate  (figs.  17  and 
18)  present  several  new  features.  One  of  these  makes  it  necessary 
to  modify  the  statement1  that  solutions  of  non-hydrated  salts  have 
practically  the  same  absorption  as  that  of  a  layer  of  water  equal  in 
depth  to  the  water  in  the  solution.  This  is  nearly  true  for  potassium 
chloride,  but  not  so  for  the  two  salts  ammonium  bromide  and  sodium 
nitrate.  In  all  four  of  the  curves  expressing  the  relations  between  the 
transmissions  of  the  solutions  and  of  the  solvent,  as  we  approach  and 
pass  through  the  centers  of  the  bands,  the  solution  is  much  less  trans- 
parent than  the  pure  solvent.  This  is  exactly  opposite  to  the  effect 
noted  in  the  case  of  hydrated  salts.  Regions  of  the  spectrum  which, 
in  the  case  of  hydrated  salts,  showed  that  the  solution  was  as  much  as 
30  to  40  per  cent  more  transparent  than  the  solvent,  show  in  the  cases 
of  both  ammonium  bromide  and  sodium  nitrate,  that  the  solution  is 
30  to  40  per  cent  less  transparent.  In  both  of  these  cases,  however, 
a  slightly  deeper  layer  of  solution  was  used  than  in  the  case  of  strongly 
hydrated  salts.  The  1/z  band  for  ammonium  bromide  (fig.  15)  shows 
that  the  solution  is  less  transparent  than  the  solvent.  It  will  be 
recalled  that  with  calcium  chloride  (figs.  11  and  12)  atypical  hydrated 
salt,  the  solution  had  the  greater  transparency  for  every  wave-length 
of  light  investigated. 

There  seems  to  be  fairly  good  evidence  that  there  is  a  shift  in  some 
of  the  absorption  bands.  If  there  is  any  shift  of  the  solution  band  for 
a  hydrated  salt,  it  is  towards  the  red.  The  I/JL  bands  for  magnesium 
chloride  (fig.  1),  magnesium  nitrate  (fig.  13),  and  zinc  nitrate  (fig.  9) 
are  displaced  slightly  towards  the  red,  and  the  centers  of  both  bands 
for  sodium  nitrate  (figs.  17  and  18)  seem  to  be  shoved  slightly  towards 
the  violet  end  of  the  spectrum.  There  does  not  seem  to  be  any  justi- 
fication for  stating  that  there  is  a  shift  in  either  direction  for  any 
absorption  band  of  any  of  the  other  solutions. 

The  question  arises,  why  are  concentrated  solutions  of  hydrated 
salts  more  transparent  than  a  layer  of  water  equal  in  depth  to  the 
water  in  the  solution?  The  answer  involves,  in  our  opinion,  the  solvate 
theory  of  solution.  From  earlier  work2  done  in  this  laboratory,  it  was 
calculated  that  in  a  solution  of  magnesium  chloride  of  the  concentra- 
tion 2.3  normal,  about  65  per  cent  of  the  water  present  is  at  ordinary 

^hys.  Zeit.,  15.  447  (1914). 
*Carnegie  Inst.  Wash.  Pub.  No.  60. 


BY   MEANS   OF   THE   RADIOMICROMETER. 


37 


temperatures  combined  with  the  dissolved  substance.     Only  35  per 
cent  of  all  the  water  present  is,  therefore,  acting  as  solvent  water. 

All  of  the  solutions  which  we  studied  were  of  much  greater  concen- 
tration than  the  above.  Since  the  total  amount  of  water  combined 
with  the  dissolved  substance  increases  with  the  concentration  of  the 
solution,  the  percentage  of  combined  water  was  greater  than  65  in  the 
solutions  with  which  we  worked.  It  has  been  pointed  out  that  only  4 


\ 


\ 


MgtNOOj  3.80N 
Depth  20 mm. 


1.05  1.10 

FIG.  13. — Wave-lengths. 


1.20 


30 


20 


Mg(N03)t  3.68 N 
Depth  10.9mm. 


1.10 


1.20 


1.25  1.30 

FIG.  14.— Wave-lengths. 


1.35 


38  STUDY   OF   ABSORPTION   SPECTRA 

of  the  14  solution-curves  given  for  hydrated  salts  show  clean-cut  evi- 
dence of  a  shift  toward  the  red.  This  would,  of  course,  cause  the  solution 
to  be  less  opaque  as  we  pass  into  and  through  the  centers  of  the  bands, 
and  more  opaque  as  we  pass  out  of  the  bands  and  reach  the  points  of 
greatest  transmission.  It  is  difficult  to  explain  why  the  IM  solution 
bands  for  magnesium  nitrate  and  zinc  nitrate  (figs.  9  and  13)  are  more 
opaque  at  the  centers  of  the  bands,  especially  as  there  is  fairly  good  evi- 
dence that  a  shift  towards  the  red  has  taken  place.  The  nitrates  have  a 
widely  different  molecular  volume  from  any  of  the  other  salts  studied. 
Their  molecular  volumes  are  very  large.  This  may  in  some  way  account 
for  the  peculiarities  in  the  transmission  curves  of  the  nitrates.  The 
sulphates  have  small  molecular  volumes  and  in  some  manner  produce 
a  concentration  in  the  solvent.  The  transmissions  at  the  centers  of 
the  1/z  bands  for  solutions  of  zinc  and  magnesium  sulphates  are  much 
higher  than  for  any  of  the  other  salts  studied  by  us  (figs.  5  and  7). 
Solutions  of  sodium  nitrate  show  that  the  nitrates  probably  have  a 
freer  vibrating  system  than  even  the  pure  solvent.  Solutions  of  this 
salt  show  at  the  centers  of  both  bands  greater  absorption,  and  at  the 
tops  of  the  bands  higher  transmission  than  the  pure  solvent.  More 
attention  should  be  given,  in  work  of  this  character,  to  the  volumes 
and  masses  of  the  vibrating  systems,  which  undoubtedly  vary  widely 
with  the  addition  of  different  salts. 

It  has  been  seen  that  in  only  4  out  of  14  curves  for  the  hydrated 
salts  is  there  any  evidence  of  a  shift  of  the  solution  curve  towards  the 
red.  In  our  opinion  the  chief  factor  which  makes  the  aqueous  solution 
more  transparent  than  a  depth  of  water  equal  to  the  wrater  in  the  solu- 
tion is  that  hydration  exists  in  the  solution.  It  is  quite  certain  now 
that  in  the  solutions  which  we  studied,  much  more  than  half  of  the 
water  was  combined  with  the  dissolved  substance.  It  seems  almost  a 
necessity  that  this  would  alter  the  vibrational  frequency  of  the  absorb- 
ing electrons  or  systems.  The  character  of  the  transmission  curves 
seems  to  justify  the  conclusion  that  water  of  hydration  has  less  power 
to  absorb  light  than  pure  uncombined  water.  We  can  conceive  of  no 
other  rational  explanation  which  will  interpret  satisfactorily  our  results 
for  hydrated  salts.  That  water  of  hydration  is  less  opaque  to  light 
than  free  water  seems,  from  our  work,  fairly  certain.  In  this  way  alone 
is  it  possible  for  us  to  explain  satisfactorily  the  transmission  curves 
for  hydrated  salts. 

A  question  of  importance  is  why,  for  non-hydrated  salts,  is  the 
transmission  curve  for  the  solution  always  below  the  transmission 
curve  for  the  solvent,  not  only  at  the  centers  of  the  bands,  but  for 
most  of  the  regions  of  the  spectra  which  we  studied?  This  is  exactly 
the  reverse  of  the  relations  already  pointed  out  for  the  hydrated  salts. 
Just  before  the  centers  of  the  bands  are  reached,  passing  in  the  direc- 
tions of  the  longer  wave-lengths,  the  solution  of  the  non-hydrated  salt 


BY   MEANS   OF   THE   RADIOMICROMETER. 


39 


is  as  much  as  35  per  cent  more  opaque  than  the  solvent.  At  present 
we  can  only  conjecture  as  to  the  real  cause  of  this  effect.  A  shift 
towards  the  violet  would  make  the  solution  more  opaque  at  the  centers 
of  the  bands;  and  although  there  is  slight  evidence  of  a  shift  in  that 
direction  for  the  bands  of  sodium  nitrate,  it  is  not  at  all  apparent  in 
the  other  bands.  The  curves  do  not  bring  out  the  characteristics  in 
general  which  would  be  expected  if  the  solution  bands  were  displaced 


H,0 


\ 


NH4Br4.02N 
Depth  22mm. 


.95 


1.005  1.10 

FIG.  15. — Wave-lengths. 


1.15 


1.20 


\ 


NH4Br  4.02  N 
Depth  I!  mm. 


1.25  1.30 

FIG.  16.— Wave-lengths. 


40  STUDY   OF   ABSORPTION    SPECTRA 

towards  the  violet.  Some  other  assumption  of  a  more  fundamental 
character  is  necessary  to  explain  the  phenomenon  under  consideration. 

If  the  non-hydrated  salts  decreased  the  association  of  the  solvent, 
and  thereby  produced  an  absorbing  system  freer  to  vibrate,  it  seems 
that  most  of  the  points  brought  out  by  the  transmission  curves  could 
be  satisfactorily  explained.  We  should  then  expect  the  solution  to  be 
more  opaque  than  the  solvent  at  the  centers  of  the  bands.  In  case 
there  was  no  shift  in  the  bands,  the  same  relations  should  hold  for  the 
other  wave-lengths  of  light  investigated.  The  above  suggestion  is  at 
present  to  be  accepted  only  tentatively,  as  offering  a  possible  explana- 
tion of  the  results  obtained. 

We  are  now  applying  a  method  which  we  hope  will  decide  whether 
the  above  assumption  is  or  is  not  correct,  whether  the  presence  of  a 
non-hydrated  salt  does  or  does  not  diminish  the  association  of  water. 

Frequent  reference  has  been  made  to  the  peculiar  characteristics  of 
the  transmission  curves  for  magnesium  nitrate.  It  was  desired  to  see 
whether  the  salt  itself  had  any  absorption.  Some  magnesium  nitrate 
was  prepared  nearly  anhydrous  by  repeated  crystallization  from  abso- 
lute alcohol,  and  the  salt  was  then  dried  in  a  vacuum  over  phosphorus 
pentoxide  for  several  days.  It  was  dissolved  in  acetone  and  a  1.07 
normal  solution  prepared.  This  was  studied  spectroscopically  in  the 
same  manner  as  has  already  been  outlined  for  aqueous  solutions. 

Fifteen  millimeters  of  acetone  showed  a  faint  band  at  0.9ju,  having 
less  than  10  per  cent  absorption.  Another  quite  similar  band  appeared 
at  1.03/z,  part  of  which  was  probably  due  to  water.  At  1.22ju  there  is 
a  fairly  intense  band  absorbing  about  60  per  cent  of  the  light.  The 
absorption  again  reaches  a  maximum  close  to  the  point  in  the  spectrum 
where  our  measurements  for  the  1.25/z  water-band  ended.  Using  such 
depths  of  the  solutions  that  they  contained  as  much  of  the  solvent  as 
had  been  employed  to  obtain  the  transmission  of  15  mm.  of  the  pure 
solvent,  it  was  hoped,  if  magnesium  nitrate  did  have  any  absorption 
of  its  own,  that  it  could  be  detected.  From  0.7/z  to  1.3/x,  which  covers 
most  of  the  spectrum  studied  in  this  work,  the  percentage  transmissions 
for  both  the  solution  and  the  solvent  were  very  nearly  the  same.  There 
was,  therefore,  no  indication  of  any  absorption  due  to  the  dissolved 
substance. 

Beyond  1.3/z,  as  we  approached  another  intense  band,  some  very  un- 
usual differences  were  noted.  Until  more  work  can  be  done  with  pure 
acetone  in  this  region,  and  with  anhydrous  magnesium  nitrate,  we  do 
not  feel  justified  in  drawing  any  conclusion  from  these  results. 

Anhydrous  magnesium  nitrate  was  also  dissolved  in  some  very  care- 
fully prepared  absolute  ethyl  alcohol,  the  concentration  of  the  solution 
being  0.53  normal.  It  was  desired  to  learn  whether  the  same  kind  of 
results  would  be  obtained  as  had  manifested  themselves  with  water 
as  the  solvent.  Table  14  is  for  20  mm.  of  alcohol.  It  will  be  noticed 


BY    MEANS    OF   THE    RADIOMICROMETER. 


41 


that  the  character  of  the  transmission  curve  for  alcohol  is  very  similar 
to  that  for  water. 

Such  depths  of  solutions  were  used  that  they  contained  in  the  one 
cell  21  mm.  of  alcohol  and  in  the  other  1  mm.  As  with  acetone,  the 
percentage  transmissions  for  both  the  solution  and  the  solvent  are  very 
nearly  the  same.  This  was  perhaps  to  be  expected,  since  with  water 
the  percentage  transmissions  of  both  solution  and  solvent  are  very 


1\ 


\ 


NaN03  5.69N 
Depth  22.2mrr 


1.05  1.10 

FIG.  17.— Wave-lengths. 


x 

\\ 

"*°\\ 

A 

\ 

\ 

Vjs~ 

^^^^ 

^\ 

\ 

V  — 

~~^^__^ 

H20             -^^ 

S 

NaN035.69N 
Depth  1  I.I  mm. 

"S 

1.10                     1.15                         1.20                        1.25                         1.30                         1.35                        1.40 
FIG.  18.  —  Wave-lengths. 

42 


STUDY    OF   ABSORPTION    SPECTRA 


nearly  identical.  The  solution  whose  percentage  transmissions  are 
given  in  table  14,  along  with  those  of  the  solvent,  is  not  nearly  so 
strongly  alcoholated  as  it  would  be  hydrated  at  the  same  concentration 
if  water  was  used  as  the  solvent.  Neither  the  work  on  magnesium 
nitrate  in  acetone,  nor  in  absolute  alcohol,  presents  any  evidence  that 
the  salt  has  any  absorption  for  the  wave-lengths  of  light  in  question. 
Jones  and  Anderson1  and  Jones  and  Strong,2  in  mapping  and  studying 
the  absorption  spectra  of  solutions,  found  abundant  evidence  for  the 
existence  of  "solvent  bands."  In  mixtures  of  certain  solvents,  say 
alcohol  and  water,  neodymium  salts  showed  simultaneously  upon  the 
plate  two  separate  and  distinct  sets  of  absorption  bands,  where  only 

TABLE  14. — Magnesium  nitrate  in  alcohol. 


X 

C2HSOH 

Mg(N03)2 
0.53  N.  in 
C2HSOH. 

X 

CzHjOH 

Mg(NO3)2 
0.53  N.  in 
C2H5OH. 

709 

\  98.0 

98  0 

1118 

79.0 

81.0 

709 

}   98.0 

98.0 

1132 

81.0 

81.8 

490 

92.3 

93.0 

1147 

80.0 

77.3 

907 

88.3 

88.7 

1162 

75.0 

65.8 

918 

79.6 

80.5 

1177 

62.4 

49.3 

934 

81.5 

81.4 

1195 

47.7 

34.0 

946 

83.2 

83.1 

1210 

32.6 

20.8 

957 

85.0 

85.3 

1227 

19.8 

16.3 

970 

86.3 

87.3 

1235 

15.5 

13.5 

983 

85.5 

86.6 

1242 

13.5 

13.1 

995 

82.0 

85.2 

1250 

13.5 

14.0 

1008 

f  79.5 

83.6 

1258 

15.3 

16.7 

1008 

1  79.0 

82.5 

1266 

18.2 

19.6 

1014 

76.8 

80.3 

1272 

21.6 

27.3 

1022 

75.0 

78.0 

;  1289 

29.2 

36.0 

1027 

73.8 

76.3 

1303 

36.7 

43.4 

1035 

72.5 

74.7 

:  1318 

43.3 

50.0 

1042 

71.5 

73.5 

1331 

47.0 

53.6 

1048 

72.6 

73.5 

1345 

49.7 

54.9 

1060 

73.5 

75.3 

1373 

49.2 

41.7 

1073 

75.8 

76.4 

1399 

35.7 

21.5 

1088 

76.0 

77.4 

1426 

15.6 

1103 

78.0 

78.0 

one  would  appear  in  each  solvent.  They  ascribed  the  one  band 
to  the  hydrate  formed  with  the  neodymium  salt  and  water,  and  the 
other  to  the  alcoholate  formed  by  the  neodymium  salt.  The  "water" 
band  and  the  "alcohol"  band  appeared  to  be  of  equal  intensity  when 
7  or  8  per  cent  of  water  was  present  in  the  alcohol. 

A  large  number  of  other  solvents  were  used,  and  "solvent"  bands 
were  found  for  each  of  them.  They  could  even  distinguish  between 
a  given  alcohol  and  its  isomer,  by  means  of  the  absorption  spectra  of 
a  neodymium  salt  dissolved  in  them.  This  was  regarded  as  good 


Carnegie  Inst.  Wash.  Pub.  No.  110. 


Ibid.,  Pubs.  Nos.  130  and  160. 


BY   MEANS    OF   THE    RADIOMICROMETER.  43 

evidence  that  the  neodymium  salts  were  solvated.  The  photographic 
results  were  hardly  more  than  qualitative.  It  was  desired  to  study 
these  "solvent"  bands  quantitatively. 

The  photographic  work  was  done  with  a  grating  which  gave  consid- 
erable dispersion.  It  was  hardly  to  be  expected  that  the  prism  in  the 
prism  spectroscope  which  we  used  would  disperse  the  light  sufficiently 
to  enable  us  to  study  the  two  "solvent"  bands.  We  could,  however, 
by  means  of  the  prism  spectroscope  and  radiomicrometer,  learn  some- 
thing of  the  difference  in  position  and  intensities  of  the  "solvent" 
bands,  and  at  least  gain  an  idea  as  to  how  to  proceed  when,  in  the  near 
future,  we  shall  study  this  problem  in  detail  with  a  large  new  grating 
spectroscope  now  completed,  and  with  a  still  more  improved  radiomi- 
crometer. 

Seven  solutions  were  prepared,  each  containing  neodymium  chloride 
of  the  concentration  0.141  normal.  The  two  solvents  were  water  and 
absolute  ethyl  alcohol.  As  recorded  in  table  15,  the  solvents  contained 
the  following  amounts  of  water:  100,  50,  25,  13,  9,  3,  and  0  per  cent. 

The  neodymium  chloride  was  of  course  dehydrated.  This  was 
effected  by  heating  for  a  long  time  at  about  150°  in  a  current  of  dry 
hydrochloric-acid  gas.  All  the  solutions  used  were  perfectly  clear. 
The  depth  of  layer  employed  in  determining  the  percentage  transmis- 
sions for  each  solution  wras  10  mm.  The  results  recorded  in  the  table 
enable  us  to  study  three  neodymium  bands  for  each  solution.  It  is 
evident  that  the  greatest  changes  in  the  percentage  transmissions  take 
place  in  those  solutions  which  contain  more  than  90  per  cent  alcohol. 
The  bands  for  only  three  of  the  solutions  have  been  plotted  (fig.  19). 
They  are  the  bands  for  neodymium  chloride  in  100  per  cent  water,  in 
97  per  cent  alcohol,  and  100  per  cent  alcohol.  The  heavy  curve  nearest 
the  violet  end  of  the  spectrum  is  for  100  per  cent  water.  The  dotted 
curve  is  for  97  per  cent  alcohol,  while  the  curve  represented  by  dashes 
is  for  100  per  cent  alcohol. 

It  is  evident  that  the  spectrum  does  not  show  any  indications  of  a 
hydrate  and  an  alcoholate  band,  both  bands  appearing  as  one.  This 
was  anticipated,  since  the  dispersive  power  of  the  glass  prism  is  so 
slight.  Near  the  centers  of  the  bands,  on  both  the  descending  and 
ascending  arms  and  at  the  center  of  the  bands,  the  wave-lengths  of 
light  recorded  in  the  table  are  those  which  would  be  obtained  by  turning 
the  calibrated  head  of  the  spectrometer  one  division.  This  increased 
or  decreased  the  wave-lengths  of  light  through  this  region  by  about 
40  A.  u.,  according  to  the  direction  in  which  the  drum-head  was  turned. 
Increasing  the  wave-length  to  this  extent  would  make  it  impossible  to 
detect  the  two  "solvent"  bands.  Accordingly,  the  drum-head  was 
turned  through  increments  of  about  5  A.  u.  and  the  transmission  deter- 
mined. This  was  done  for  the  mixed  solvents  for  all  three  bands  given 
in  the  curves,  and  also  for  a  few  of  the  bands  which  showed  best  the 


44 


STUDY   OF   ABSORPTION    SPECTRA 


"solvent"  bands  when  the  photographic  method  was  employed. 
Working  in  this  way,  even  with  the  small  dispersion  of  the  glass  prism, 
fairly  definite  indications  were  obtained  of  the  existence  of  "hydrate" 
and  "alcoholate"  bands.  It  will,  however,  be  necessary  to  use  a  spec- 
troscope with  higher  resolving  power  to  bring  out  more  clearly  the  exist- 
ence of  "solvent"  bands.  Such  a  piece  of  apparatus  is  now  completed. 
The  curves  in  figure  19  present  many  interesting  features.  The 
intensity  of  absorption  for  neodymium  chloride  in  100  per  cent  water 
is  much  greater  than  that  for  neodymium  chloride  in  100  per  cent 
alcohol.  It  can  be  stated,  in  general,  that  the  intensity  of  the  absorp- 


.90 


FIG.  19. — Wave-lengths. 


tion  increases  as  the  water-content  increases.  Neodymium  chloride 
in  water  has  a  band  whose  center  is  at  A  =  0.739/*,  and  here  the  percent- 
age transmission  is  about  29.  For  neodymium  chloride  in  alcohol  it 
is  nearly  50  per  cent  for  the  above  wave-length.  The  transmission  of 
neodymium  chloride  in  water  for  this  wave-length  of  light  is  therefore 
about  40  per  cent  less  than  the  transmission  for  this  same  salt  in  pure 
alcohol.  The  ratio  of  intensities  of  the  hydrate  and  of  the  alcoholate 
bands  is  about  the  same  at  0.8/z,  where  the  center  of  the  second  band 
for  neodymium  chloride  in  water  is  located.  The  much  less  intense 
neodymium  water-band,  whose  center  is  at  X  =  0.875^,  is  barely  30  per 


BY   MEANS    OF   THE    RADIOMICROMETER. 


45 


cent  more  intense  than  that  for  neodymium  chloride  in  pure  alcohol. 
On  the  ascending  arms  of  the  curves  exactly  the  opposite  effects  are 
noted.  This  indicates  a  displacement  of  the  alcoholate  band  to  the 
region  of  longer  wave-lengths.  The  displacement  is  very  large  in  the 
case  of  the  first  band.  Neodymium  chloride  in  100  per  cent  water  has 
a  band  whose  center  is  at  X  =  0.739/i.  When  the  solvent  is  100  per 

TABLE  15. — Results  for  neodymium  chloride  in  mixed  solvents.     Concentration 
0.141  normal. 


X 

100  p.  ct. 
H20 

50  p.  ct. 
H20 

25  p.  ct. 
H2O 

13  p.  ct. 
H2O 

9  p.  ct. 
H20 

3  p.  ct. 
H20 

0  p.  ct. 
H20 

702 

/  97.7 

98.0 

98.0 

97.6 

100.0 

100.0 

100.0 

702 

\  98.1 

98.0 

98.0 

98.0 

98.0 

98.0 

98.0 

709 

95.9 

96.3 

96.2 

98.0 

98.1 

98.1 

98.2 

717 

96.3 

94.8 

95.0 

96.2 

98.2 

98.2 

98.3 

725 

89.6 

91.8 

93.6 

91.3 

95.3 

96.5 

96.8 

732 

51.7 

52.3 

55.6 

55.2 

62.8 

74.6 

88.2 

736 

42.3 

41.3 

42.2 

40.0 

44.4 

50.7 

68.2 

739 

28.8 

32.0 

31.8 

30.5 

34.4 

40.9 

49.3 

743 

f  29.5 

31.8 

22.3 

29.5 

32.8 

36.1 

41.3 

743 

\  29.9     32.0 

31.9 

29.7 

32.2 

34.8 

40.7 

749 

f  40.0  !   43.0 

40.0 

37.5 

35.3 

31.2 

35.7 

749 

I  41.8 

43.0 

39.7 

37.7 

35.9 

32.4 

35.2 

753 

/  58.0 

59.7 

58.7 

54.3 

52.7 

45.2 

41.8 

753 

1  58.7 

59.5 

59.4 

55.4 

53.5 

45.2 

41.6 

756 

78.3 

75.0 

76.0 

75.6 

73.7 

65.7 

62.1 

764 

91.1 

91.0 

89.6 

91.3 

91.6 

89.3 

85.2 

773 

94.1 

92.5 

93.5 

96.6 

97.6 

95.6 

93.6 

780 

84.5 

83.6 

85.5 

85.7 

89.3 

91.6 

93.9 

789 

58.4 

56.7 

57.7 

61.5 

64.6 

74.0 

88.3 

794 

32.2 

32.2 

32.6 

34.1 

38.0 

53.1 

70.3 

799 

25.0 

21.8 

21.8 

21.6 

21.3 

27.8 

48.6 

803 

f  31.6 

32.3 

31.3 

29.0 

29.0 

24.5 

29.0 

803 

1  32.7 

32.0 

32.7 

29.2 

28.8 

24.0 

30.1 

807 

f  50.8 

49.0 

50.5 

48.2 

45.6 

41.5 

32.4 

807 

\  51.3 

49.3 

51.4 

48.2 

46.5 

40.9 

33.1 

813 

68.7 

69  5 

68.2 

65.7 

64.5 

57.3 

49.6 

817 

80.3 

79.8 

80.2 

77.5 

75.8 

68.0 

64.2 

823 

85.6 

86.3 

86.5 

84.7 

82.7 

79.4 

76.2 

828 

90.8 

90.5 

90.7 

90.2 

90.7 

87.3 

84.8 

837 

94.2 

92.5 

95.6 

93.7 

95.3 

95.7 

94.7 

847 

93.6 

91.6 

95.3 

94.8 

95.5 

95.9 

95.1 

858 

90.5 

90.0 

92.0 

92.3 

94.3 

95.5 

96.1 

869 

71.6 

70.6 

73.7 

74.5 

77.4 

84.1 

91.0 

875 

60.7 

60.8 

60.8 

62.3 

65.5 

72.4 

84.1 

875 

60.3 

61.2 

61.2 

62.8 

65.5 

72.3 

82.7 

880 

66.3 

66.7 

66.3 

65.7 

65.6 

66.3 

73.6 

880 

66.4 

66.0 

66.3 

65.6 

66.0 

66.3 

73.0 

885 

76.5 

75.8 

77.6 

74.8 

74.0 

70.3 

68.1 

885 

77.5 

76.0 

77.0 

75.0 

73.4 

71.9 

68.8 

890     82.3 

82.1 

83.5 

81.3 

80.8 

77.6 

71.3 

896 

83.6 

82.4 

80.7 

76.7 

903     83.6 

82^9 

84*5 

S3.Q 

82.5 

80.6 

907     86.2 

84.5 

82.4 

918 

90.3 

86^3 

87^0 

85.6 

85!  1 

84.3 

84.7 

934 

89.5 

82.6 

89.2 

946 

85.6 

89^5 

90~3 

91.7 

46  STUDY    OF   ABSORPTION    SPECTRA 

cent  alcohol,  the  alcoholate  band  has  its  center  at  X  =  0.749//,  the  dis- 
placement here  being  about  100  A.  u.  For  the  second  band  the  dis- 
placement is  only  about  40  A.  u.,  while  for  the  third  it  is  over  100  A.  u. 
When  the  percentage  of  alcohol  is  below  90,  the  bands  for  the  mixed 
solvents  are  only  slightly  displaced  towards  the  red,  the  greatest  change 
taking  place  between  90  and  100  per  cent  alcohol. 

Similar  experiments  were  carried  out  with  anhydrous  neodymium 
chloride  dissolved  in  pure  glycerol  and  in  water.  The  salt  was  also 
dissolved  in  mixtures  of  the  two  solvents.  The  glycerol  bands  were 
very  slightly,  if  at  all,  displaced  towards  the  longer  wave-lengths. 

It  gives  us  pleasure  to  express  our  thanks  to  Professor  H.  A.  Pfund 
and  to  Professor  John  A.  Anderson,  for  many  valuable  suggestions 
during  the  progress  of  this  work. 


CHAPTER  II. 

CONDUCTIVITIES,  TEMPERATURE  COEFFICIENTS  OF  CONDUCTIVITY, 

DISSOCIATIONS,  AND  CONSTANTS  OF  CERTAIN  ORGANIC 

ACIDS  IN  AQUEOUS  SOLUTIONS. 

BY  LESLIE  D.  SMITH. 

A  part  of  Dr.  Smith's  work  with  the  organic  acids  has  already 
appeared  in  Publication  No.  170  of  the  Carnegie  Institution  of  Wash- 
ington. The  remainder  of  his  work  is  included  in  this  chapter. 

INTRODUCTION. 

This  investigation  is  a  continuation  of  work  which  has  been  in 
progress  during  the  past  12  years  on  the  conductivity,  temperature 
coefficients  of  conductivity,  and  dissociation  of  electrolytes  between  0° 
and  65°.  The  work  was  undertaken  in  connection  with  the  solvate 
theory  of  solution,  which  was  proposed  by  Jones  about  15  years  ago. 

The  first  work  was  done  by  Jones  and  West1  on  certain  salts,  and 
extended  over  the  temperature  range  0°  to  35°. 

The  second  investigation  was  by  Jones  and  Jacobson2  on  a  number 
of  salts,  over  the  same  range  in  temperature. 

The  work  of  White  and  Jones3  was  on  the  conductivity,  dissociation, 
and  dissociation  constants  of  a  number  of  organic  acids  from  0°  to  35°. 

The  fourth  investigation  was  by  Clover  and  Jones,4  using  organic 
acids  and  salts.  The  conductivities  were  measured  from  35°  to  80°. 
This  was  the  first  piece  of  work  done  in  this  laboratory  over  the  higher 
range  in  temperature. 

The  second  piece  of  work  of  White  and  Jones5  was  on  the  conduc- 
tivity, dissociation,  and  dissociation  constants  of  a  number  of  organic 
acids  from  0°  to  35°. 

The  second  investigation  of  Jones  and  West6  was  on  a  number  of 
salts  between  35°  and  65°. 

"The  seventh  investigation  in  this  field  was  by  Wightman  and  Jones,7 
on  the  conductivity,  dissociation,  and  dissociation  constants  of  organic 
acids  between  0°  and  35°. 

The  eighth,  by  Hosford  and  Jones,8  had  to  do  with  the  conductivities, 
temperature  coefficients  of  conductivity,  and  dissociation  of  certain 
electrolytes  from  0°  to  35°. 

The  ninth,  by  Winston  and  Jones,9  dealt  with  the  conductivities, 
temperature  coefficients  of  conductivity,  and  dissociation  of  certain 
salts  from  0°  to  35°.  Miss  Winston  independently  worked  out  an 
interesting  theory  of  induction  in  solution. 

^mer.  Chem.  Journ.,  34,  357  (1905).     4Ibid.,  43,  187  (1910).     ''Ibid.,  46,  56  (1911). 
2Ibid.,  40,  355  (1908).  *Ibid.,  44,  159  (1910).     *Ibid.,  46,  240  (1911). 

3Ibid.,  42,  520  (1909).  'Ibid.,  44,  508  (1910).     'Ibid.,  46,  368  (1911). 

47 


48  CONDUCTIVITIES    AND    DISSOCIATIONS 

The  second  investigation  by  Wightman  and  Jones1  had  to  do  with  the 
conductivity  and  dissociation  of  certain  organic  acids  from  35°  to  65°. 

The  eleventh  investigation  in  this  field  was  by  Springer  and  Jones,2 
on  the  conductivity,  dissociation,  and  dissociation  constants  of  a  large 
number  of  organic  acids  between  0°  and  65°. 

The  twelfth  piece  of  work  was  done  by  Howard  and  Jones3  on  the 
conductivity,  temperature  coefficients  of  conductivity,  and  dissociation 
of  certain  electrolytes  between  35°  and  65°. 

The  thirteenth  research  was  carried  out  by  Shaeffer  and  Jones4  on 
the  conductivity,  temperature  coefficients  of  conductivity  and  dissocia- 
tion of  certain  salts.  They  also  studied  the  effect  of  hydration  and 
hydrolysis.  A  number  of  salts  which  had  presented  apparent  abnor- 
malities in  the  earlier  work  were  studied. 

All  of  the  above  work,  and  that  which  was  done  by  Smith  over 
the  temperature  range  35°  to  65°,  has  been  discussed  in  Publication 
No.  170  of  the  Carnegie  Institution  of  Washington. 

The  foregoing  brief  survey  of  the  investigations  in  this  laboratory 
is  all  that  is  necessary  here,  since  the  papers  by  White  and  Jones5  and 
by  Wightman  and  Jones,6  already  referred  to,  contain  a  careful  review 
of  the  work  previously  done  on  the  conductivity  of  organic  acids  in 
aqueous  solution. 

PURPOSE  OF  THE  INVESTIGATION. 

The  object  of  this  investigation  was  to  secure  more  data  concerning 
the  dissociation  of  organic  acids  over  a  wide  range  in  temperature  and 
dilution,  to  improve  methods,  to  test  the  work  already  done,  and  to 
discover  new  relations  between  the  additional  data  obtained. 

EXPERIMENTAL. 

The  investigation  here  discussed  extended  over  about  two  years. 
The  first  year  was  devoted  to  a  study  of  the  conductivity  and  disso- 
ciation of  certain  organic  acids  at  35°,  50°,  and  65°.  The  acids  used 
were  some  of  those  that  had  been  studied  at  the  lower  temperatures 
by  other  workers.  The  readings  at  35°  were  repeated,  and  the  agree- 
ments were  very  satisfactory.  Where  the  difference  was  appreciable 
the  work  was  repeated.  If  the  difference  was  still  too  great  to  be  due 
to  unavoidable  experimental  error,  the  readings  at  the  lower  tempera- 
tures were  taken  again  and  a  complete  set  of  new  data  secured.  The 
agreements  were  very  close  in  practically  every  case,  although  the 
readings  were  taken  by  different  investigators. 

The  work  during  the  last  year  was  devoted  to  the  study  of  several 
acids  at  the  lower  temperatures,  0°,  12.5°,  25°,  and  35°. 

iAmer.  Chem.  Journ.,  48,  320  (1912).          3Ibid.,  48,  501  (1912).  6Ibid.,  44,  156  (1910). 

*Ibid.,  48,  411  (1912).  *Ibid.,  49,  207  (1913).  6Ibid.,  46,    56  (1911). 


OF  CERTAIN  ORGANIC  ACIDS  IN  AQUEOUS  SOLUTIONS.   49 

The  acids  which  had  already  been  investigated  by  White  and  Jones 
at  the  lower  temperatures,  and  which  were  reinvestigated  by  the  author 
were,  hippuric,  gallic,  picric,  and  crotonic.  In  every  case  very  close 
agreements  were  found,  and  the  values  at  35°  agreed  to  within  the 
limits  of  experimental  error.  The  temperatures  were  the  same  as  those 
used  by  the  former  investigators. 

REAGENTS. 

The  water  used  was  purified  in  the  same  manner  as  that  employed 
in  the  earlier  investigations.  This  method  was  described  in  detail  by 
Jones  and  Mackay.1 

The  acids  were  obtained  from  Kahlbaum  and  Schuchardt.  The 
general  method  of  purification  was  recrystallization  from  conductivity 
water.  The  butyric  acids  —  normal  and  iso  —  were  purified  by  distilla- 
tion in  a  vacuum;  acetaminobenzoic  acid  was  recrystallized  from  ether.2 
Whenever  possible  the  acids  were  carefully  dried  in  a  vacuum  desiccator 
containing  sulphuric  acid.  If  practicable,  the  melting-points  of  the 
acids  were  taken  as  one  criterion  of  purity.  The  "mother"  solution, 
i.  e.,  the  one  with  the  greatest  concentration,  was  made  up  whenever 
possible  by  direct  weighing,  and  then  titrated  against  standard  alkali. 
The  acids  which  are  liquids  were  made  up  directly  by  titration.  The 
modification  of  this  method  for  the  higher  temperatures  was  described 
in  detail  in  the  paper  by  Wightman  and  Jones.3  All  flasks  and  burettes 
used  in  this  investigation  were  calibrated  by  the  method  of  Morse  and 
Blalock,4  and  also  by  weight  of  the  contained  water. 

The  sodium  salts  of  the  organic  acids  were  used  to  determine  the  /z«, 
values  of  the  acids,  and  were  prepared  by  titration.  It  had  been  found 
that  the  sodium  salts  of  organic  acids,  in  general,  are  completely  disso- 
ciated at  a  dilution  of  F  =  2048.  Therefore,  a  solution  of  the  acid  at 
this  dilution  was  just  neutralized  by  sodium  hydroxide,  using  phenol- 
phthalein  as  the  indicator,  and  its  conductivity  determined  in  the 
usual  way. 

APPARATUS. 

The  cells  used  in  this  investigation  were  the  same  as  those  employed 
by  Jones  and  Wightman.  In  their  paper  they  state:5 

"The  cells  resembled  those  used  by  Jones  and  Bingham,6  with  platinum- 
plate  electrodes  attached  to  glass  tubes  containing  mercury,  the  tubes  being 
sealed  into  ground  glass  stoppers.  As  many  as  eight  cells  were  employed  with 
constants  ranging  from  about  10  to  330  in  Siemens  units.  A  cell  of  special 
type,7  having  a  very  low  constant,  was  used  for  obtaining  the  conductivity  of 
the  water.  In  order  to  get  a  sharp  reading  in  the  cells,  the  electrodes  were 
covered  with  a  fine  coating  of  platinum  black,  in  the  usual  manner." 


r.  Chem.  Journ.,  19,  91  (1897).  *>Ibid.,  46,  62  (1911). 

2Carnegie  Inst.  Wash.  Pub.  No.  170.  «Ibid.,  34,  493  (1905). 

3Amer.  Chem.  Journ.,  48,  320  (1912).  ''Ibid.,  45,  282  (1911). 
*Ibid.,  16,  479  (1894). 


50  CONDUCTIVITIES   AND   DISSOCIATIONS 

The  thermometers  employed  were  carefully  compared  with  a  stan- 
dard Reichsanstalt  thermometer.  The  standard  thermometer  was  also 
calibrated  at  the  United  States  Bureau  of  Standards. 

The  resistance  box  which  was  used  throughout  this  entire  investiga- 
tion had  also  been  calibrated  by  the  Bureau  of  Standards.  A  very 
fine  slide  wire  bridge  was  employed,  the  wire  being  wrapped  around  a 
marble  drum.  On  this  bridge  the  investigator  could  easily  read  to 
fractions  of  a  millimeter.  The  entire  scale  was  5  meters  long,  and  read- 
ings could  be  made  to  tenths  of  a  millimeter. 

The  cells  were  kept  at  a  constant  temperature  in  thermostats.  As  this 
work  was  carried  on  simultaneously  with  that  of  Dr.  Wightman,  and 
partly  using  his  apparatus,  reference  is  made  to  his  description  of  the 
system  of  thermal  regulation.  In  the  paper  by  Wightman  and  Jones1 
the  system  of  regulation  for  temperatures  0°  to  35°  is  fully  described. 
They  say: 

"Three  thermostats  were  employed  to  keep  the  cells  at  constant  tempera- 
ture; one  for  0°  similar  to  that  described  by  Jones  and  Jacobson;2  one  for  15° 
(in  this  case  for  12.5°)  and  25°,  a  galvanized  tub  containing  20  to  30  liters  of 
water,  and  in  the  bottom  of  which  was  placed  a  lead  (in  this  investigation, 
copper)  coil  through  which  cold  water  was  passed  under  constant  pressure; 
a  third  for  35°,  differing  from  the  latter  only  in  not  having  a  coil  in  the  bottom. 
They  were  both  kept  constantly  stirred  by  propellers  driven  by  a  hot-air 
engine.  In  this  way  it  was  possible  to  keep  the  temperature  constant  to 
within  0.02°." 

The  work  at  the  more  elevated  temperatures  was  carried  out  in 
apparatus  similar  to  that  used  at  the  lower  temperature,  with  the 
difference  that  the  thermostats  were  covered  with  asbestos  boards 
saturated  with  paraffin.  The  thermometer  and  stirrer  passed  through 
small  openings  in  the  cover,  and  the  cells  were  placed  in  the  bath 
through  openings  just  large  enough  to  admit  them.  These  holes  were 
covered  by  pieces  of  the  same  material  as  the  cover.  The  evaporation 
of  the  water  in  the  bath  was  in  this  way  reduced  to  a  minimum,  and 
the  confined  vapors  maintained  the  cells,  both  the  exposed  portion  and 
the  part  immersed  in  the  bath,  at  the  same  temperature.  Thermo- 
regulators  of  the  same  type  as  are  used  throughout  this  laboratory  were 
employed. 

There  has  been  much  less  work  done  here  at  higher  than  at  lower 
temperatures;  therefore,  a  more  detailed  description  of  the  various 
difficulties  encountered,  and  the  methods  devised  to  overcome  them, 
will  be  given. 

The  only  work  that  had  been  previously  carried  out  at  the  higher 
temperatures  is  that  of  Clover  and  Jones3  and  of  West  and  Jones.4 

lAmer.  Chem.  Journ.,  46,  56  (1911).  3Ibid.,  43,  187  (1910). 

*Ibid.,  40,  355  (1908).  *IUd.,  44,  508  (1911). 


OF   CERTAIN   ORGANIC   ACIDS   IN   AQUEOUS   SOLUTIONS.        51 

PROCEDURE. 

The  cell  constants  were  taken  by  the  method  described  by  White  and 
Jones.1  When  working  at  low  temperatures  they  were  determined  as 
often  as  once  a  month,  and  still  more  frequently  when  working  at  high 
temperatures.  The  experimental  difficulties  are  much  more  numerous 
at  the  higher  than  at  the  lower  temperatures.  There  is  a  greater  and 
more  frequent  variation  in  the  cell  constants.  The  solubility  of  the 
glass  is  often  too  large  to  be  negligible.  A  few  short  quotations  from 
the  work  of  Clover  and  Jones,  referred  to  above,  will  give  the  results 
of  their  investigations  in  the  best  possible  forms. 

CELL   CONSTANTS. 

"It  has  developed  that  a  strain  is  brought  about  by  the  high  temperatures 
which  may  result  in  a  change  either  in  the  distance  of  the  electrode  plates  from 
each  other  or  in  the  surface  of  the  plates.  Since  such  a  variation  had  not 
previously  been  observed  in  work  covering  a  range  of  0°  to  35°,  it  was  thought 
that  the  changes  might  be  reduced  by  maintaining  the  cells  at  a  temperature 
which  was  about  a  mean  of  those  employed  in  the  experimental  work.  Accord- 
ingly, when  not  in  use,  the  cells  were  filled  with  pure  water  and  placed  in  a 
bath  which  was  maintained  continuously  at  a  temperature  of  45°  to  50°." 

They  performed  an  experiment  to  test  the  value  of  this  method. 

"The  measurements  were  first  carefully  made  at  35°  and  then  duplicated; 
then  the  regular  systematic  procedure  was  gone  through  at  65°,  and  after  this 
the  readings  at  35°  were  again  made.  If  the  results  found  the  second  time 
should  agree  with  those  first  obtained  at  35°,  this  would  be  strong  evidence 
that  the  method  was  reliable.  It  was  found  in  some  cases  that  the  second 
reading  differed  slightly  from  the  first.  In  other  cases  there  was  no  difference, 
and  the  change  appeared  to  be  independent  of  the  cell  used,  or  the  concen- 
tration of  the  solution.  It  was  further  found  that  on  standing  for  a  con- 
siderable length  of  time  at  35°  (2  or  3  hours),  the  reading  slowly  changed  back 
in  all  cases  to  the  original  value.  It  is  difficult  to  see  what  other  causes  can  be 
assigned  to  these  results  than  a  temporary  change  in  the  cell  constant  during 
the  heating  at  65°." 

In  conductivity  work  carried  out  at  ordinary  temperatures,  experi- 
ments have  shown  that  the  error  introduced  by  the  solubility  of 
the  glass  is  negligible.  However,  at  50°  the  error  from  this  same  source 
at  a  dilution  of  one  liter  is  the  largest  of  all  the  ordinary  experimental 
errors.  The  glass  is  still  more  soluble  at  65°,  as  would  be  expected. 
Clover  and  Jones  point  out  that  at  80°  the  conductivity  of  pure  water 
is  increased  tenfold  on  standing  in  the  cells  for  a  couple  of  hours.  The 
cells  used  in  their  research  were  made  of  hard  glass.  Obviously,  the 
amount  of  glass  dissolved  from  the  cells  depends  largely  upon  the  nature 
of  the  glass  of  which  the  cells  are  made.  It  varies  considerably  with 
the  different  cells.  The  attempt  to  introduce  directly  a  correction 
factor  for  the  solubility  of  glass  was  therefore  abandoned.  This  source 
of  error  was  overcome  in  another  way. 

.  Chem.  Journ.,  42,  520  (1909). 


52  CONDUCTIVITIES   AND   DISSOCIATIONS 

It  was  found  that  after  the  cells  had  been  heated  with  water  for 
several  hours,  the  amount  of  glass  dissolved  gradually  decreased  and 
finally  became  practically  nothing.  After  this  treatment,  as  the  cells 
were  kept  in  a  bath  at  45°  and  65°  and  the  water  in  them  changed  once 
a  day,  the  solubility  of  the  glass  at  65°  was  always  practically  zero  and 
therefore  negligible.  If  a  cell  was  removed  from  the  bath  and  allowed 
to  stand  filled  with  water  at  room  temperature  for  any  appreciable 
length  of  time,  it  was  found  on  heating  that  glass  was  again  dissolved. 

In  this  investigation  the  above  precautions  were  all  observed.  In 
addition,  at  the  start  of  the  work  the  glass  vessels  were  heated  with 
a  dilute  solution  of  caustic  soda,  then  boiled  with  dilute  hydrochloric 
acid,  and  then  for  some  time  with  conductivity  water.  The  solubility 
of  the  glass  was  shown  by  tests  to  have  been  reduced  to  a  negligible 
quantity  by  this  treatment.  The  cells  were  then  kept  filled  with  con- 
ductivity water  which  was  changed  every  day,  and  were  maintained 
at  a  temperature  approximately  50°,  and  thus  the  solubility  of  the  glass 
was  negligible.  This  was  proved  experimentally.  A  cell  was  filled 
with  conductivity  water  at  35°  and  the  conductivity  read ;  then  it  was 
heated  and  readings  were  made  at  both  50°  and  65°.  The  solution  was 
cooled  to  35°  and  the  conductivity  again  measured.  Since  the  two 
readings  at  35°  coincided,  there  was  no  appreciable  quantity  of  glass 
dissolved  during  the  process. 

This  was  not  the  only  beneficial  result  obtained  from  keeping  the 
cells  always  at  an  elevated  temperature.  The  expansion  of  glass  and 
of  platinum  is  not  large  enough  to  affect  appreciably  the  cell  constants, 
as  would  be  the  case  if  the  temperature  varied  over  a  wider  range. 

It  has  already  been  stated  that  it  takes  several  hours  for  a  cell  to 
acquire  its  normal  condition  at  the  lower  temperature,  with  respect  to 
the  electrode-plates  and  the  glass  rods  when  suddenly  cooled.  This 
fact  was  utilized  in  the  determination  of  the  cell  constants  at  the  higher 
temperatures.  The  cells  were  kept  at  about  50°,  and  when  the  cell 
constants  were  to  be  taken  the  cells  were  filled  with  the  solution  of 
potassium  chloride  and  they  were  heated  to  50°  for  about  an  hour. 
The  cells  containing  the  solution  were  then  cooled  down  and  the  con- 
ductivity read  immediately  after  a  constant  temperature  had  been 
reached.  The  readings  were  made  at  25°,  where  the  conductivity  of 
the  solutions  of  potassium  chloride  is  accurately  known. 

The  method  of  White  and  Jones  was  used  for  taking  the  cell  con- 
stants over  the  entire  range  of  temperature  from  0°  to  65°. 

Two  objects  were  accomplished  in  all  the  work  at  higher  tempera- 
tures, by  filling  the  cells  nearly  to  the  top  with  the  solutions.  First, 
evaporation  and  the  consequent  change  in  concentration  were  reduced 
to  a  minimum.  Second,  carbon  dioxide  and  other  gases  from  the  air 
were  prevented  from  being  present  in  sufficient  quantity  to  affect 
the  results. 


OF   CERTAIN   ORGANIC   ACIDS   IN   AQUEOUS   SOLUTIONS.        53 

The  method  of  preparing  the  solutions  was  that  devised  by  West 
and  Jones : 

"Since  we  worked  over  a  range  of  temperature  of  only  30°,  we  found  it 
convenient  to  prepare  the  solutions  at  the  intermediate  temperature,  50°,  and 
then  to  use  the  solutions  at  the  three  temperatures,  35°,  50°,  and  65°.  But 
since  the  volume  of  a  solution  varies  with  the  temperature,  it  was,  of  course, 
necessary  to  apply  a  correction  at  35°  and  65°  to  the  volume  of  solutions 
made  up  to  50°. 

"When  a  standard  solution  is  cooled  from  50°  to  35°  there  is  a  contraction 
in  volume  and  a  consequent  increase  in  the  concentration  of  the  solution. 
The  value  of  n,  for  any  solution  would,  therefore,  be  slightly  too  large.  The 
value  of  nv  as  found  must  be  multiplied  by  the  factor  0.994  for  results  at  35° 
when  the  solutions  were  made  up  at  50°.  The  correction  factor  for  solutions 
made  up  at  50°  and  used  at  65°  is  1.0076.  The  coefficient  of  expansion  for 
distilled  water  is  somewhat  less  than  that  for  an  aqueous  solution.  However, 
the  difference  in  the  coefficients  for  water  and  for  that  of  our  most  concentrated 
solution  is  so  small  that  it  is  negligible.  By  making  use  of  the  above  correction 
it  was  necessary  to  prepare  only  one  set  of  solutions  for  each  salt;  and,  conse- 
quently, much  pure  material  and  time  were  saved." 

When  Wightman,  Springer,  and  Smith  began  the  investigations  at  the 
higher  temperatures  the  two  problems  presenting  themselves  were : 

(1)  In  making  up  the  solutions  at  50°  the  glass  of  the  flasks  would 
dissolve,  because  the  flasks  were  not  kept  at  50°  in  a  bath,  but,  when 
not  in  use,  were  allowed  to  cool  down  to  room  temperature. 

(2)  The  solution,  even  if  heated  to  50°,  would  cool  off  when  poured 
into  the  burette,  and  still  further  when  draining  into  the  flask;  there- 
fore, instead  of  drawing  off  50  c.c.  of  the  solution  at  50°,  50  c.c.  would 
be  drawn  off  at  a  temperature  of  50°  —  x°,  or  more  than  50  c.c.  at  50°. 

This  complex  source  of  error  was  overcome  by  the  above-named 
workers,  all  of  whom  started  simultaneously  the  present  line  of  work 
at  elevated  temperatures.  The  results  are  given  in  the  paper  by 
Springer  and  Jones  i1 

"A  simple  device  did  away  with  both  of  these  sources  of  error.  At  50°, 
988.07  grams  of  distilled  water  have  a  volume  of  1000  c.c.  Our  liter  flask 
was  weighed  at  20°.  Then  988.07  grams  of  distilled  water  at  20°  were  intro- 
duced— air  displacement  being  taken  into  account.  The  flask  was  marked  at 
the  bottom  of  the  meniscus.  This  flask,  filled  to  the  mark  with  water  at  20° 
(room  temperature),  will  contain  a  liter  at  50°.  Therefore,  all  the  mother 
solutions  could  be  made  up  for  50°  work  at  20°,  and  the  solubility  of  the  glass, 
which  is  noticeable  only  at  higher  temperatures,  is  thus  made  negligible.  This 
mother  solution  will  be,  let  us  say,  normal  at  50°.  At  20°  it  is  stronger  than 
normal.  If  we  draw  out  100  c.c.  at  20°  and  dilute  it  to  200  c.c.  at  20°,  it  will 
also  be  stronger  than  half -normal.  But  heat  this  latter  solution  to  50°  and  it 
will  again  attain  its  required  normality — assuming  that  the  coefficient  of 
expansion  of  water  is  the  same  as  that  of  dilute  solutions.  And  all  solutions 
worked  with  were  eighth-normal,  or  more  dilute.  This  permits  the  titration 
of  all  solutions  at  room  temperature,  which  obviates  the  second  of  the  above 
difficulties." 

.  Chem.  Journ.,  48,  417  (1912). 


54 


CONDUCTIVITIES   AND    DISSOCIATIONS 


This  solution  of  the  above-mentioned  problems  was  worked  out  by 
Jones  and  Springer  and  Smith  jointly,  and  with  the  same  apparatus. 
We  used  a  500  c.c.  or  a  200  c.c.  flask  calibrated  in  this  way  for  50°  at 
20°,  in  which  to  make  up  our  mother  solutions.  In  the  higher  temper- 
ature work  it  was  noted  that  bubbles  of  air  formed  on  the  electrodes, 
especially  at  50°  and  60°.  These  were  removed  by  carefully  shaking 
the  cells  before  making  the  readings. 

VALUES. 

The  values  of  the  limiting  conductivities  of  the  monobasic  acids  were 
found  directly  from  the  /*=*,  values  of  their  sodium  salts.  It  has  been 
shown  by  experiment  that  the  conductivities  of  sodium  salts,  made  up 
by  the  titration  method,  which,  as  already  stated,  was  used  in  this 
work,  agree  with  the  conductivities  of  solutions  which  were  made  up 
directly  from  the  dry,  solid,  sodium  salt.  The  ju«  values  of  the  acid 


P        •(>       12       14       16       18       20       22       24      26       28      30 

FIG.  20. — Limiting  conductivities. 

were  obtained  from  those  of  the  sodium  salts  by  a  very  simple  method. 
It  consists  in  subtracting  the  constant  of  the  sodium  ion  from  the 
limiting  conductivity  value  for  the  sodium  salt,  and  then  adding  to  this 
value  the  constant  for  the  hydrogen  ion.  This  may  be  readily  accom- 
plished in  actual  work  by  using  this  equation:  ju*,  acid  =  /j.x  HC1 +/*<*> 
Na  salt  of  acid  —  ^«,  NaCl. 

To  calculate  the  value  of  ju«,  for  hydrochloric  acid  and  sodium 
chloride  at  the  various  temperatures  at  which  the  work  was  done,  the 
equations  of  White  and  Jones  were  used : 

For  sodium  chloride,     /*»  =  63.04  +  2.04  t  -  0.00823  t2 
For  hydrochloric  acid,  /*»  =  245.4  +  6.06  t  -  0.00776  t2 
Table  16  gives  all  the  values  of  the  acids  with  which  Smith  worked. 
The  /i<»  values  of  dibasic  acids  could  not  be  determined  by  this  method. 
The  sodium  salts  of  dibasic  acids  do  not  yield  a  maximum  value  of 
conductivity  at  dilutions  at  which  work  could  be  done.     Comparison 
of  the  ju°°  values  of  the  acids  which  Jones  and  several  of  his  co-workers 
have  studied,  shows  that  those  acids  with  the  largest  number  of  atoms 


OF   CERTAIN   ORGANIC   ACIDS   IN   AQUEOUS   SOLUTIONS.        55 

in  the  anions  have  the  smallest  /*«  values.  A  curve  in  which  the 
ordinates  represent  the  /*«>  values  of  the  acids,  and  abscissas  the 
number  of  the  atoms,  was  plotted  for  all  these  acids.1  By  placing  the 
dibasic  acids  on  this  curve  according  to  the  number  of  atoms  present, 
their  values  were  obtained.  (See  fig.  20.) 

The  values  for  the  acids  marked  with  an  asterisk  (*)  in  table  16 
were  determined  by  other  workers  in  this  laboratory,  and  can  be  found 
in  publication  No.  170  of  the  Carnegie  Institution  of  Washington, 
pages  91  and  92.  It  was  shown  that  acids  having  the  same  number 

TABLK  16. — Limiting  conductivities  of  the  acids. 


Acid. 

0° 

12.5° 

25° 

35° 

50° 

65° 

Butyric*  
Isobutyric*  

404 
403 

473.3 
437  29 

540.3 

468  7 

Malic  

221.5 

286  2 

352 

400 

Pyrotartaric*  
Racemic*  
Hippuric  
Citric* 

219 

280 

345 

397 
398 
392 
392 

468 
468.2 
446.4 
464  5 

533 
534.9 
499.8 
528  5 

220  2 

285 

350 

398 

222 

286 

352 

402 

475  1 

544 

Maleic*  

402 

475 

544 

Fumaric*  

402 

475 

544 

Itaconic*  

400 

471 

537.5 

Citraconic*  
Mesaconic*  
m-Chlorobenzoic  
p-Chlorobenzoic  
o-Bromobenzoic  
m-Bromobenzoic  
m-Hydroxybenzoic  
m-Acetoxybenzoic  

222^4 
222 
223 

222.8 
219.7 
220.4 
220 

285^3 

285.2 
286.5 
285 
285.8 
284 
284  5 

350^5 
353.5 
352.5 
352.6 
350.8 
349.1 
349  5 

400 
400 
398.5 
400 
401 
401.6 
397.3 
397 
397  5 

471 
471 

537.5 
537.5 

m-Sulphobenzoic  
Picric  
Gallic  
Aminobenzenesulphonic  
o-Toluic*  

220 
207 
220 
221 

284.5 
281.6 

'285' 

349.5 
331 
348 
350 

397.5 
380 
396 
398 
397 

451  '.  6 
459.1 

470.1 

5CKL8 
514.3 

536 

m-Toluic*  
p-Toluic*  
Mandelic  

218 

284 

343  5 

397 
397 
397 
391 

470.4 
469.1 
475.4 

541.9 

538.8 
539.9 

of  atoms  in  the  molecule  have  practically  the  same  /*»  value.  From 
these  values  and  the  others  given  in  the  table  just  referred  to  the  curve 
was,  therefore,  plotted.  On  this  curve  it  was  possible  to  find  the  values 
of  ju*>  for  the  dibasic  acid  by  means  of  the  graphic  method  previously 
described.  The  curve  shows  that  at  the  higher  temperatures  the  method 
is  far  less  accurate,  since  the  values  which  are  determined  directly, 
using  the  sodium  salt,  fall  less  frequently  on  a  definite  curve  at  the 
higher  temperatures. 


'Carnegie  Inst.  Wash.  Pub.  No.  170. 


56  CONDUCTIVITIES   AND    DISSOCIATIONS 

The  results  of  the  conductivity  measurements  are  tabulated  in  the 
following  pages.  The  molecular  conductivities  (/x,)  are  expressed  in 
Siemens  units.  Temperature  coefficients  are  expressed  in  both  conduc- 
tivity units  and  percentages.  The  former  are  obtained  by  the  equation 

(Hn  —  MI)/(*I  ~~  0  =  conductivity  units ; 
the  latter  by  the  equation 

(Uti—  M<)/M*(^i~ t)=per  cent  temperature  coefficient1 
where  t=  the  lower  temperature,  ti  the  next  higher  temperature,  & 
and  jutl  the  conductivities  of  the  same  solution  measured  at  tempera- 
tures t  and  t\,  respectively. 

Obviously  the  percentage  dissociation  may  be  expressed  as 

a=/*,/M°° 

The  dissociation  constants  are  obtained  from  Ostwald's  dilution  law 

a2/(l-a)7  =  X 

This  is  found  to  hold  for  the  weaker  organic  acids.  In  the  tables 
where  no  values  are  given  for  KX104,  the  acids  were  so  strong  that 
they  did  not  obey  the  law. 

The  following  acids  were  tested,  but  were  found  to  be  either  too 
slightly  soluble  to  work  with,  or  to  undergo  decomposition : 

Nitrocinnamic  acid  (insoluble). 

jS-Naphtholdisulphonic  acid  (insoluble). 

Terephthalic  acid  (insoluble). 

Aminosalicylic  acid  (insoluble). 

Aminophenolsulphonic  acid  (decomposed). 

p-  and  o-nitro-  and  dinitro-phenols  were  studied,  but  the  conductivi- 
ties found  were  so  small  that  the  results  were  not  reliable,  since  the  per- 
centage error  due  to  unavoidable  experimental  errors  was  relatively  large. 

In  table  18  omitted  values  have  been  determined  by  other  investi- 
gators, especially  those  whose  values  are  given  at  only  the  higher 
temperatures. 

It  has  been  shown  by  the  work  of  Ostwald,  Jones,  and  others  in 
this  field,  that  the  Ostwald  dilution  law  does  not  hold  for  strongly 
dissociated  acids.  The  " constants"  of  acids  of  this  class  have  been 
omitted.  In  the  tables  that  were  secured  in  the  experimental  work, 
irregular  constants  are  frequently  given.  These  are  of  value  only  in 
determining  the  relative  strengths  of  such  acids. 

It  had  been  previously  observed  that  as  a  rule  the  percentage  dis- 
sociations of  the  organic  acids  decreased  with  rise  in  temperature,  but 
that  the  values  of  percentage  dissociation  of  certain  acids  increased  with 
rise  in  temperature.  In  this  work,  however,  it  was  found  that  in  each 
case  the  increase  has  been  at  a  diminishing  rate,  and  further  study 
showed  that  the  maximum  dissociation  was  reached  in  all  cases  except 
citric  acid. 

'Zeit.  phys.  Chem.,  2,  561  (1888). 


OF   CERTAIN   ORGANIC   ACIDS   IN   AQUEOUS   SOLUTIONS.        57 


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58 


CONDUCTIVITIES    AND    DISSOCIATIONS 


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OF    CERTAIN    ORGANIC    ACIDS    IN    AQUEOUS    SOLUTIONS.         59 


In  the  work  of  Springer1  the  maximum  dissociation  was  reached  in 
every  case  studied  by  him,  except  for  the  amino  acids,  such  as  sulpha- 
nilic,  metanilic,  naphthionic,  and  dinitroaminophenol.  Smith  did  not 
study  acids  of  this  class;  but  certain  groups  of  those  which  he  has 
investigated  show  an  increase  in  dissociation  over  a  wide  range  in 
temperature.  The  acids  that  show  this  property  to  the  greatest  extent 
are  those  containing  the  hydroxyl  (OH)  group.  These  acids  have  per- 
centage dissociation  values  which  increase,  attain  a  maximum,  and 
then  decrease. 

TABLE  18. — Dissociation  constants. 


Acid. 

0° 

12.5° 

25° 

35° 

50° 

65° 

m-Chlorobenzoic  
p-Chlorobenzoic  
o-Bromobenzoic*  
m-Bromobenzoic  
m-Hydroxybenzoic  
m-Acetoxybenzoic  
o-Sulphobenzoic*  
m-Sulphobenzoic*  
Picric*  

1.41 
0.55 

1^33 
0.65 
1.19 

1.52 
0.62 

1.49 
0.67 
1.29 

1.57 
0.78 

1.52 
0.68 
1.41 

0.38 

1.54 
0.68 

1.55 

0.69 
1.32 

6.39 

1.25 
0.55 
0.44 
4.2 
2.5 
0.140 
0.114 
5.00 

0.42 

i.oo 

0.52 
0.39 
3.7 

0.131 
0.121 

6  40 

6  88 
0.46 
0.35 
3.5 

0.118 
0  122 

2  23 

6  188 

8  8 
1.54 
31.5 

6.4 

Gallic  

0.36 

Aminobenzenesulphonic*  
o-Toluic*  

m-Toluic*  
p-Toluic*  
Mandelic*  
Sebacic  
Butyric*  
Isobutyric*  
Malic  
Pyrotartaric  
Racemic* 

2.8 

4.28 
0.88 

2.7 

4.67 
0.85 

2.6 

4'si 

0.83 

Hippuric  
Citric* 

2.11 

2.20 

2.27 

2.23 

2.32 

6  '623 

9~8 
1.55 
32.9 
7.1 

Aconitic* 

0.208 

11^3 
1.5 

35.8 

7.8 

Crotonic  

0.193 

0.206 

0.210 

Maleic*  

Fumaric*  

Itaconic*  

Citraconic*  
Mesaconic*  

*  Value  found  by  some  other  investigator;    see  Publication    of   the   Carnegie   Institution 
of  Washington,  No.  170. 

The  temperature  coefficients  of  conductivity  expressed  in  conduc- 
tivity units,  for  all  of  the  above  acids  show  an  increase  with  the  dilution 
of  the  solution  and  a  decrease  with  rise  in  temperature.  As  would  be 
expected,  the  rate  of  increase  or  decrease  varies  greatly.  There  seems 
to  be  a  relation  between  the  strength  of  the  acid  and  rate  of  increase 
or  decrease  of  the  temperature  coefficients  of  conductivity  with  rise  in 
temperature.  The  stronger  the  acid  the  less  rapid  is  the  change. 
It  should  be  noted  that  strong  acids  are  so  largely  dissociated  at  all 

1Amer.  Chem.  Journ.,  48,  411,  (1912). 


60 


CONDUCTIVITIES   AND   DISSOCIATIONS 


TABLE  19. 


dilutions  that  a  change  in  the  dilution  would  change  the  coefficients 
very  slightly. 

The  temperature  coefficients  of  conductivity  of  acids,  even  of  mineral 
acids  to  a  certain  extent,  decrease  with  rise  in  temperature.  The 
literature  shows  also  that  the  mineral  acids  are  more  or  less  hydrated 
in  solution.  It  is  evident  that  an  acid  which  is  either  unhydrated  or 
only  slightly  hydrated  (as  is  the  case  with  most  of  the  organic  acids 
studied  in  this  investigation)  would  show  a  rapid  decrease  in  the 
temperature  coefficients  of  conductivity  with  rise  in  temperature; 
since  there  would  be  no  complex  hy- 
drated ion  to  lose  water,  and,  thus, 
from  this  cause,  the  conductivity 
could  not  be  increased.  This  point 
is  illustrated  in  table  19,  which  gives 
the  temperature  coefficients  of  certain 
weak,  unhydrated  acids  at  F  =  1024. 

Table  20  from  the  work  of  Jones 
and  Wightman,1  gives  similar  values  for  some  of  the  hydrated  acids; 
two  weaker  acids — racemic  and  citric — show  the  same  decrease  in 
the  coefficients  as  the  strong  acids.  This  indicates  that  the  strength 
of  the  acid  has  very  little  to  do  with  the  gradual  decrease  in 
temperature  coefficients,  but  that  this  decrease  is  caused  by  hydra- 
tion;  and  the  above  relation  applies  to  the  strong  acids  only  because 
they  are  the  acids  which  are  the  most  strongly  hydrated.  This  same 

TABLE  20. 


Acid. 

35°-50° 

50°-65° 

Butyric  
Isobutyric  
Pyrotartaric  

0.019 
0.015 
0.12 

0.013 
0.009 
0.11 

Acid. 

Water  of 
crystalli- 
zation 

0°-15° 

15°-25° 

25°-35° 

Percent- 
age 
decrease. 

H2O 

Meconic  

3 

10.06 

9.94 

9.18 

8.8 

Benzenesulphonic  

1 

5.26 

5.05 

5.06 

3.8 

p-Toluenesulphonic  

4 

4.92 

4.84 

4.67 

5.3 

1,  2,  4-Nitrotoluenesulphonic  

2.5 

4.74 

4.67 

4.35 

8.2 

Racemic  KQ°  =  9.1  

1 

3.76 

3.56 

3.43 

8.0 

Citric  K0°  =  6.85  

1 

3.64 

3.59 

3.38 

7.1 

relation  holds  in  this  work.  The  temperature  coefficients  of  conduc- 
tivity of  the  acids  are  larger  for  stronger  acids.  This  is  shown  by 
table  21,  in  which  the  acids  are  arranged  according  to  their  strengths. 
All  the  work  done  in  the  Chemical  Laboratory  of  the  Johns  Hop- 
kins University  on  the  conductivity  of  organic  acids  has  shown  that 
the  temperature  coefficients  of  conductivity  expressed  in  percentage 
decrease  slightly  with  increasing  dilution  and  more  rapidly  with  rise 
in  temperature.  The  decrease  due  to  rise  in  temperature  manifests 


.  Chem.  Journ.,  46,  56  (1911). 


OF  CERTAIN  ORGANIC  ACIDS  IN  AQUEOUS  SOLUTIONS.    61 


itself  in  every  case  studied,  but  a  large  number  of  exceptions  present 
themselves  in  the  decrease  due  to  increasing  dilution.  No  doubt  this 
is  the  result  of  hydrolysis. 

TABLE  21. 
[Temperature  coefficients,  35°-50°;  F=1024.] 


Acid. 

Cond.  units. 

K 

0  316 

0  114 

0  54 

0  208 

1  62 

1  55 

Mandelic  

1.83 

4.2 

5  12 

100  per  cent  diss 

In  all  the  acids  that  Smith  studied,  the  percentage  temperature 
coefficients  of  conductivity  are  small  and  are  of  the  same  order  of 
magnitude.  The  decrease  is  regular  as  the  temperature  rises.  This 
indicates  that  some  constant  factor  has  influenced  the  solution,1  such 
as  the  viscosity  of  the  medium. 

The  isomeric  acids  have  been  the  subject  of  a  large  amount  of  investi- 
gation. Ostwald,2  Walker,3  and  Derrick4  have  worked  on  the  correla- 
tion of  ionization  and  structure  for  such  acids,  and  on  the  effect  of 
substitution.  This  work  has  confirmed  the  following  statement  of 
Springer  and  Jones : 

"In  the  case  of  isomeric  acids  containing  two  carboxyl  groups,  the  acid  with 
these  groups  in  the  'ortho'  or  the  'cis'  position  is  always  the  stronger." 


o-Toluic  35°, 
m-Toluic  35°, 
p-Toluic  35°, 
m-Bromobenzoic  35C 


K  =  1.25 
K  =  0.55 
K  =  0.44 
K  =  1.55 


Maleic  (cis)  35°, 
Fumaric  (trans)  35°, 
Citraconic  (cis)  35°, 
Mesaconic  (trans)  35C 


K=  154 
K  =  11.3 

K  =  35.8 
K  =  7.8 


o-Bromobenzoic  35°,  too  strong  to  give  a  constant. 


Euler5  showed  that  the  conductivity  of  organic  acids  is  a  parabolic 
function  of  the  temperature.  This  was  also  brought  out  by  the  work 
of  White  and  Jones.  Wightman  and  Jones,  and  Springer  and  Jones 
plotted  curves,  using  conductivities  as  ordinates  and  temperatures  as 
abscissas,  and  calculated  the  conductivity  of  a  number  of  acids  from 
Euler's  formula:  = 


This  was  always  found  to  hold.     It  holds  also  in  this  work. 

The  dissociation  of  a  large  number  of  organic  acids  decreases  with 
rise  in  temperature.  As  mentioned  above,  in  certain  cases  the  per- 
centage dissociation  apparently  attains  a  maximum  at  one  of  the  inter- 
mediate temperatures,  usually  at  25°  or  35°.  This  has  been  observed 
and  noted  in  the  publications  of  Euler,6  Schaller,7  White  and  Jones,8 


1  Jones  and  West;  Amer.  Chem.  Journ.,  34, 418  (1905). 
2Zeit.  phys.  Chem.,  3,  170  (1889). 
3Journ.  Chem.  Soc.,  61,  605  (1892);  67   147  (1895). 
4Journ.  Amer.  Chem.  Soc.,  33, 1881  (1911) ; 34, 74  (1912). 


5Zeit.phys.  Chem., 21,  257  (1896). 

*Ibid.,  21,  247  (1896). 

7 Ibid.,  25,  497  (1898). 

8Amer.  Chem.  Journ.,  44, 196  (1910). 


62 


CONDUCTIVITIES   AND   DISSOCIATIONS 


and  Wightman  and  Jones,1  but  thus  far  no  one  has  offered  a  satisfactory 
explanation.  The  problem  is  not  simple,  there  being  many  factors, 
any  one  of  which  might  greatly  influence  the  dissociation.  As  is  well 
known,  decrease  in  the  association  of  the  solvent  and  in  its  dielectric 
constant  would  greatly  affect  the  conductivity.  The  effect  of  rise  in 
temperature  has  been  shown  to  diminish  both  the  association  and  the 
dielectric  constant  of  the  solvent.  The  nature  of  the  dissolved  sub- 
stance itself  plays,  of  course,  no  small  role  in  the  effect  on  the  conduc- 
tivity of  its  solution. 

Table  22  shows  agreement  in  results  in  the  cases  of  a  few  acids 
studied.     The  examples  were  chosen  at  random  and  are  typical. 

TABLE  22. 


Acid. 

V. 

White  and 
Jones. 

Smith  and 
Jones. 

Ostwald. 

Maleic,  35°  

f       32 
1024 
[  2048 

198.8 
384.6 
400.8 

197.6 
384.97 
400.10 

Gallic,  35°  

(       64 
1024 
1  2048 

19.36 
71.18 
96.7 

19.46 
71.20 
96.82 

Hippuric,  25"  

f     128 
I  1024 

55.16 
133.96 

54.3 
131.1 

m-Hydroxybenzoic,  25°  .... 

f       32 
512 
(  1024 



18.38 
68.24 
93.15 

18.18 
67.90 
91.63 

Gallic,  25°  

{       64 
512 
1024 
[  2048 

16.90 
48.33 
62.50 
85.02 

16.92 
48.31 
62.56 
85.10 

16.83 

47.74 
66.53 

Malic,  25°  

f     128 
512 
1    1024 
1  2048 

78.46 
132.70 
172.50 
211.16 

71.52 
128.1 
166.6 
213.0 

SUMMARY. 

Most  of  the  relations  previously  established  by  Jones  and  his  co- 
workers  have  been  confirmed  in  this  investigation.  The  mere  state- 
ment of  some  of  these  relations  is  sufficient,  since  a  glance  at  the  data 
alone  will  show  that  they  hold. 

1.  The  temperature  coefficients  of  conductivity  increase  rapidly  with 
dilution,  and  decrease  rapidly  with  rise  in  temperature  for  the  weak 
organic  acids.     However,  if  the  acids  are  hydrated  the  temperature 
coefficients  of  conductivity  are  larger,  and  the  rate  of  increase  or  decrease 
just  mentioned  is  much  slower.     The  stronger  the  organic  acid  the 
larger  these  units. 

2.  The  percentage   temperature   coefficients   of   conductivity   are 
small,  are  of  the  same  order  of  magnitude,  and  decrease  with  rise  in 
temperature. 


1Amer.  Chem.  Journ.,  46,  56  (1911). 


OF   CERTAIN   ORGANIC   ACIDS   IN   AQUEOUS   SOLUTIONS.        63 

3.  The  conductivity  of  most  organic  acids  is  a  parabolic  function  of 
the  temperature,  as  has  been  pointed  out  by  Euler.     His  formula 


has  been  shown  to  hold  in  every  case. 

4.  The  relative  strength  of  organic  acids  is  not  dependent  upon  the 
temperature. 

5.  "There  is  no  general  statement  possible  concerning  the  change  in 
dissociation  of  the  organic  acids  with  change  in  temperature.     Maxima 
occur  with  several  between  25°  and  35°,  while  in  other  cases  maxima 
are  indicated  at  slightly  higher  temperatures  than  those  at  which 
measurements  were  made.     The  dissociation  of  several  acids  decreases 
regularly  from  0°." 

6.  Isomeric  acids  do  not  behave  similarly  with  regard  to  change  in 
their  dissociation. 

7.  Strong  organic  acids,  as  is  well  known,  do  not  obey  the  Ostwald 
dilution  law. 

8.  The  "ortho"  and  "cis"  forms  of  isomeric  acids  are  the  stronger, 
i.  e.,  have  the  larger  constants. 

9.  The  migration  velocities  of  the  anions  are  functions  of  the  num- 
ber of  atoms  that  are  present  in  them.     As  has  already  been  pointed 
out  in  this  paper,  this  fact  was  made  use  of  in  determining  the  AU  values 
of  dibasic  acids. 

10.  Most  dibasic  acids  dissociate  like  monobasic  acids. 

11.  "  The  behavior  of  the  organic  acids  with  respect  to  the  change  in 
their  dissociation  with  the  temperature,  is  not  in  accord  with  the 
hypothesis  of  Thomson-Nernst,  which  connects  dissociating  power  and 
dielectric  constants;  or  at  least  the  influence  of  some  other  unknown 
factor  is  suggested." 


CHAPTER  III. 

A  PRELIMINARY    STUDY    OF    THE    CONDUCTIVITY    OF    CERTAIN 
ORGANIC  ACIDS  IN  ETHYL  ALCOHOL  AT  15°,  25°,  AND  35°. 

BY  E.  P.  WlGHTMAN  AND   J.  B.  WlESEL. 

Considerable  work  has  been  done,  in  the  Chemical  Laboratory  of 
the  Johns  Hopkins  University,  especially  within  the  last  five  years, 
on  the  conductivity  and  dissociation  of  organic  acids  in  water  as  a 
solvent,  over  a  fairly  wide  range  of  temperature  and  dilution;  and  a 
number  of  interesting  relationships  have  been  pointed  out.1  Since,  up  to 
the  present,  almost  nothing  has  been  done  with  these  acids  in  absolute 
alcohol,  we  decided  to  extend  our  investigations  on  the  organic  acids 
into  this  field. 

HISTORICAL. 

A  few  rather  crude  measurements  of  the  conductivity  of  hydrochloric 
acid  and  some  inorganic  salts  had  been  made  in  alcohol-water  mixtures 
as  early  as  1882.  The  first  work  which  need  be  mentioned  is  that  of 
Kablukoff2  in  1889.  He  determined  the  conductivity  of  hydrochloric 
acid  in  alcohol  and  some  other  solvents,  and  in  mixtures  of  alcohol  and 
water,  and  showed  that  molecular  conductivity  increased  markedly  with 
decrease  in  amount  of  alcohol  and  increased  slowly  with  the  dilution. 

Wakeman,3  in  1893,  measured  the  conductivities  of  some  organic 
acids,  as  well  as  of  hydrochloric  acid,  potassium  chloride,  and  sodium 
chloride,  in  mixtures  of  alcohol  and  water  ranging  from  pure  water  to 
50  per  cent  alcohol.  Using  Lentz's4  values  for  the  relative  migration 
velocities  of  hydrogen,  sodium,  and  potassium  in  water  and  mixtures 
of  water  and  alcohol,  and  employing  the  method  of  Ostwald5  for 
determining  the  ju«>  values  for  the  organic  acids  from  their  sodium 
salts,  he  calculated  the  dissociations  of  these  acids.  He  also  calculated 
dissociation  constants  by  means  of  Ostwald's  dilution  law.  His  dis- 
sociation values  decrease  slowly  with  increase  in  alcohol.  The  con- 
stants decrease  much  more  rapidly  for  the  same  increase  in  alcohol. 
Wakeman  plotted  curves  (fig.  21)  with  molecular  conductivities  as 
ordinates  and  percentage  alcohol  as  abscissa,  and  showed  that  when 
they  were  extended  beyond  50  per  cent  alcohol  in  the  direction  of  100 
per  cent  alcohol,  the  conductivity  probably  approached  zero  as  a  limit. 
He  concluded  that  dissociation  in  the  mixtures  is  much  less  than  would 

Publication  170,  Part  II,  Carnegie  Inst.  Wash.;  see  also  Amer.  Chem.  Journ.,  44,  156  (1910)  ; 

46,  56  (1911);  48,  320,  411  (1912);  50,  1  (1913). 
2Zeit.  phys.  Chem.,  4,  429  (1889). 
3Ibid.,  11,  49  (1893). 

4Mem.  de  1'Acad.  de  St.  Petersb.,  VII  series,  vol.  30,  9  (1882). 
6Zeit.  phys.  Chem.,  2,  270  (1888) ;  ibid.,  3,  170  (1889) ;  see  also,  Amer.  Chem.  Journ.,  46,  66  (1911). 

64 


CONDUCTIVITIES   OF  ORGANIC  ACIDS  IN   ETHYL   ALCOHOL.      65 


be  expected,  and  that  the  Ostwald  dilution  law  could  not  be  applied 
to  the  mixtures  containing  much  alcohol. 

In  1894  Schall1  determined  the  conductivity  of  oxalic,  dichloracetic, 
picric,  and  hydrochloric  acids  in  methyl  alcohol,  in  ethyl  alcohol, 
in  ethyl  alcohol-water  mixtures,  and  in  isobutyl  alcohol.  He  con- 
cluded from  his  results  that 
molecular  conductivity  is  much 
less  in  the  alcohols  than  in 
water,  and  that  the  acids  be- 
have very  differently  in  alcohol- 
water  mixtures  than  in  the  pure 
solvents;  some  acting  just  the 
opposite  from  what  might  be 
expected  from  their  behavior  in 
the  pure  solvents.  For  in- 
stance, picric  acid  gives  a  much 
higher,  and  the  others  much 
lower,  conductivity  values  in 
water-alcohol  mixtures  than  in 
the  pure  alcohol. 

A  careful  piece  of  work  on 
the  conductivity  of  certain 
organic  acids,  acetic,  mono- 
chloracetic,  dichloracetic,  tri- 
chloracetic,  and  succinic  acids, 
and  of  hydrochloric  acid  in 
absolute  alcohol  at  18°  was 
carried  out,  in  1894,  by  Wil- 
dermann.2  The  alcohol  was 
freed  from  aldehyde  by  treat- 
ment with  silver  nitrate,  and 
from  water  by  heating  with 
calcium  oxide.  Great  care 
was  exercised  in  protecting 
the  alcohol  from  the  air,  and 
a  special  apparatus  was  con- 
structed for  drawing  a  meas- 
ured quantity  of  alcohol  out  of 
the  supply  bottle  directly  into 
the  conductivity  cell,  which  had  a  capacity  of  about  25  c.c.  In  order 
to  make  his  minimum  points  on  the  bridge  more  distinct,  Wildermann 
employed  a  graphite  resistance  rheostat  which  he  constructed  and 
standardized  against  a  known  resistance.  He  states  that  his  con- 
ductivity cells  had  to  be  washed  with  running  water  for  about  8  to  10 


FIG.  21. — Monobromacetic  acid. 


!Zeit.  phys.  Chem.,  14,  701  (1894). 


2Ibid.,  14,  231  (1894). 


66  CONDUCTIVITIES   OF   ORGANIC   ACIDS 

days.  On  washing  with  alcohol  and  drying,  it  was  found  that  the 
alcohol  which  was  in  contact  with  the  platinum  electrodes  was  oxid- 
ized by  the  air,  even  when  this  was  thoroughly  purified  and  dried,  to 
acetic  acid,  which,  of  course,  increased  the  conductivity  of  the  alcohol. 
Therefore,  instead  of  drying  after  washing  with  water  and  draining, 
alcohol  was  introduced  so  as  first  to  wash  the  glass  walls,  and  then  the 
alcohol  was  allowed  to  cover  the  electrodes.  Some  of  this  alcohol  was 
then  drawn  off  by  means  of  a  pipette  and  fresh  alcohol  was  introduced, 
keeping  the  electrodes  continuously  covered,  until  all  the  water  had 
been  eliminated  and  the  conductivity  values  after  each  removal  and 
addition  of  fresh  alcohol  remained  unchanged.  Says  Wildermann,  "this 
cost  a  half  day  of  work  and  300  to  500  c.c.  of  good  absolute  alcohol." 
The  same  procedure  was  adopted  for  dilute  solutions  of  the  acids,  the 
strength  of  any  solution  being  determined  by  titration.  He  does  not 
give  any  tables  of  his  results  with  the  weaker  acids — acetic,  mono- 
chloracetic,  and  succinic.  He  simply  makes  the  qualitative  statement 
that  these  substances  between  volumes  10  and  160  give  a  molecular 
conductivity  which  increases  approximately  proportional  to  the  square 

root  of  the  volumes;   that  is  ^!L  is  about  equal  to  ,.  pi,  v  being  the 

greater  of  the  two  volumes. 

Wildermann  draws  the  following  conclusions: 

"(1)  Under  about  10  liters  for  dichloracetic  acid,  the  values  of  —  are 

Mr 

less  than  those  of  + —',  above  about    10  liters  the   value  of  —  is 

_     \»  Mr 

greater  than  J^l;  in  dilutions  from  800  to  2,000  liters  the  values  of 

—  become  almost  equal  to  — .     The  increase  of  —  is  therefore  con- 

M.  v  Mr 

tinuous,  not  only  in  the  concentrated  solutions,  but  also  in  the  more 

dilute  solutions  to  which  the  equation— — ^ — -v=k  may  be  applied." 

"(2)  The  same  conclusions  are  even  more  nearly  true  in  the  case  of 
/3-resorcylic  acid." 

(3)  In  the  case  of  trichloracetic  acid  there  is,  just  as  above,  an  increase 

in  —  from  20  to  300  liters,  above  which  the  —  value  becomes  almost 

Mr  Mr 

equal  to     I^L  and  then  increases  throughout  the  more  dilute  solutions. 

\v 

That  this  is  true,  even  at  a  different  temperature  from  18°,  is  shown 
in  some  later  work  by  the  author,  using  an  entirely  independent  method. 
An  explanation,  however,  had  already  been  offered  for  the  phenomenon 


IN   ETHYL   ALCOHOL.  67 

in  a  previous  paper  by  the  author.1  In  this  earlier  work  he  was  studying 
the  effect  of  the  presence  of  the  large  amount  of  the  undissociated  por- 
tion of  weakly  dissociated  acids,  upon  the  conductivity  and  dissocia- 
tion values  and  upon  the  dissociation  constants.  It  had  been  pointed 
out  by  Ostwald  that  in  aqueous  solutions  the  degree  of  dissociation, 
for  the  same  dilutions  of  trichloracetic,  dichloracetic,  monochloracetic, 
and  acetic  acids  shows  a  decrease  in  passing  from  the  trichlor  derivative 
to  the  acetic  acid  in  the  order  named.  A  like  succession  was  observed 
by  Wildermann  for  the  same  acids  in  alcohol. 

(4)  For  hydrochloric  acid  in  alcohol,  results  analogous  to  those  in 
water  were  obtained.  A  maximum  value  of  the  conductivity  was  noted. 

In  summing  up,  Wildermann  says  that  it  is  possible  to  apply  the 
Kohlrausch  method  to  the  determination  of  the  conductivity  of  strong 
organic  or  inorganic  acids  in  absolute  alcohol,  but  that  no  reliable 
results  could  be  obtained  for  such  weak  acids  as  acetic,  monochloracetic, 
and  succinic.  He  remarks  that  much  time  and  patience  on  the  part  of 
the  experimenter  are  required  to  obtain  results  that  are  at  all  reliable. 

In  a  second  investigation  by  Wildermann2  the  same  acids  as  in  the 
earlier  work  were  studied,  using  in  this  case  a  precision  galvanometer 
method  and  working  at  25°  instead  of  at  18°.  He  arrived  at  precisely 
the  same  conclusions  as  before,  except  that  he  found  the  precision 
method  susceptible  of  more  general  application  than  that  of  Kohlrausch. 

Among  those  who  have  worked  on  conductivity  in  alcohol  since 
Wildermann  are,  Zelinsky  and  Krapiwin,3  who  determined  the  conduc- 
tivity of  a  number  of  inorganic  salts  and  acid  salts  of  organic  acids ; 
Ernest  Cohen,4  who  obtained  the  conductivity  and  dissociation  of 
several  inorganic  salts  in  absolute  methyl  alcohol  at  18°,  and  who 
states  that  because  of  the  action  of  the  platinum  electrodes  upon  the 
solutions,  measurements  were  unsafe  at  higher  temperatures;  Roth,5 
whose  work  had  to  do  with  the  conductivity  of  potassium  chloride  in 
alcohol-water  mixtures;  and  others  of  minor  importance. 

Some  still  more  recent  work6  has  been  done  in  this  laboratory  with 
inorganic  salts  both  in  methyl  and  in  ethyl  alcohol  and  alcohol-water 
mixtures,  but  it  need  not  be  discussed  here.  (See  Chapter  IV.) 

EXPERIMENTAL. 

The  conductivity  apparatus  used  for  making  the  measurements 
was  similar  to  that  employed  in  previous  work  in  this  laboratory,  except 
that  on  account  of  the  high  resistances  offered  by  the  alcoholic  solutions 
of  the  acids,  it  was  necessary  to  make  use  entirely  of  the  cylindrical 
type  of  conductivity-cell.  The  method  of  obtaining  their  constants 
has  previously  been  described7  and  need  not  be  dealt  with  here. 

^Ber.  d.  chem.  Gesell.,  26,  1782-1783  (1893).  *Ibid.,  42,  209  (1903). 

2Zeit.  phys.  Chem.,  14,  247  (1894).  «Carnegie  Inst,  Wash.  Pubs.  Nos.  80  and  180. 

*Ibid.,  21,  35  (1896).  7Amer.  Chem.  Journ.,  42,  527  (1909);  44,  64 

4Ibid.,  25,  1  (1898).  (1911). 


68 


CONDUCTIVITIES   OF   ORGANIC   ACIDS 


Since  the  percentage  temperature  coefficients  of  conductivity  for 
substances  in  alcohol,  as  well  as  the  coefficient  of  expansion  of  the 
alcohol  itself,  are  so  large,  it  is  necessary  to  have  fairly  close  tempera- 
ture regulation.  This  was  secured  by  the  combination  of  a  specially 
devised  gas  regulator  and  thermo-regu- 
lator.  These  have  already  been  de- 
scribed1 in  earlier  papers. 

In  cooperation  with  Dr.  P.  B.  Davis, 
of  the  Chemical  Laboratory  of  the  Johns 
Hopkins  University,  a  new  form  of  con- 
stant-temperature bath  was  also  de- 
signed. Its  construction  can  be  seen 
from  fig.  22.  A  full  discussion  of  the 
form  finally  adopted  will  be  presented  in 
a  paper  soon  to  be  published  by  Jones, 
Davis,  and  Putnam.2  In  these  baths  the 
temperature  ordinarily  does  not  vary 
more  than  0.02°  C.,  which  is  sufficiently 
constant  for  our  purpose.  With  special 
precautions  as  to  insulation  from  changes 
in  temperatures,  and  a  further  modified 
form  of  the  thermo-regulator,  the  varia- 
tion can  be  decreased  to  a  few  thousandths 
of  a  degree.  Aside  from  the  better 
temperature  regulation  obtained  in  this 
new  form  of  thermostat-bath,  there  are 
also  one  or  two  other  advantages  derived  from  its  use.  The  apparatus 
is  of  copper  which  does  not  rust,  and  the  stirring  arrangements  and 
the  cooling  coil  are  on  the  side,  and  are  therefore  out  of  the  way. 
A  number  of  minor  improvements  were  likewise  added. 

Solutions  were  made  up  in  200  c.c.  flasks  calibrated  for  25°,  and 
the  conductivity  measurements  of  these  solutions  were  taken  at  15°, 
25°,  and  35°.  Pipettes  were  frequently  used  for  measuring  purposes 
because  of  greater  ease  in  handling.  They  were  carefully  calibrated. 
Corrections  for  the  expansion  and  contraction  of  the  alcoholic  solutions 
at  35°  and  15°,  respectively,  were  of  course  applied  to  the  conductivity 
measurements. 

The  alcohol  was  prepared  by  heating  ordinary  95  per  cent  alcohol 
for  several  days  with  fresh,  unslaked  lime  in  a  copper  tank,  provided 
with  a  ground-brass  stopper  and  reflux  condenser,  and  then  distilling 
through  a  block-tin  condenser.  The  distillate  thus  prepared  was 
reheated  with  fresh  lime  and  again  distilled,  the  first  and  last  portions 

^eit.  phya.  Chem.,  85,  519  (1913);  Journ.  Chim.  Phys.  July  1914. 
2See  Chapter  VI  of  this  monograph. 


IN   ETHYL   ALCOHOL.  69 

of  this  distillate  being  discarded.  A  few  sticks  of  sodium  hydroxide 
added  during  the  last  day  of  heating  insured  the  removal  from  the 
distillate  of  any  aldehyde  which  might  have  been  present,  and  which 
otherwise  would  have  distilled  over  with  the  alcohol.  It  is  possible, 
by  taking  proper  precautions  in  the  manner  of  handling,  to  obtain  by 
such  a  method  alcohol  having  a  specific  gravity  of  0.78506,  to  within 
the  limits  of  experimental  error,  +0.00002.  According  to  Circular 
19  of  the  Bureau  of  Standards,  alcohol  with  this  specific  gravity  has 
no  water  in  it ;  that  is,  it  is  100  per  cent  alcohol.  The  alcohol  employed 
in  the  conductivity  measurements  varied  in  specific  gravity  from  the 
value  of  0.78506  to  0.78517,  the  latter  containing  99.964  per  cent 
alcohol.  The  receiver  for  the  distillate  was  a  6-liter  Jena  glass  bottle. 
The  stopper  was  a  three-holed  paraffined  cork.  Through  one  hole 
passed  a  siphon,  through  another  an  adapter  with  a  glass  stopcock, 
and  through  the  third  a  calcium  chloride-soda  lime  tube  also  having  a 
glass  stopcock.  In  this  way  the  alcohol  was  well  protected  during 
distillation  from  impurities  in  the  air,  and  small  quantities  sufficient 
for  making  up  the  solutions  could  be  drawn  off  without  exposing  the 
main  supply.  After  weighing  out  the  quantity  of  dried  and  purified  acid 
necessary  to  make  a  solution  of  the  required  normality,  the  acid  was 
washed  off  the  watch  glass  or  out  of  the  weighing  bottle  into  a  funnel, 
and  then  into  a  200  c.c.  Jena  flask  which  had  previously  been  thoroughy 
washed  with  water,  and  then  with  some  of  the  alcohol  with  which  the 
solution  was  to  be  made  up.  The  flask  was  filled  to  the  neck  with 
alcohol  and  shaken  until  all  the  acid  had  dissolved.  It  was  finally 
hung  in  a  25°  thermostat-bath  until  temperature  equilibrium  was 
reached,  and  then  filled  to  the  mark.  In  the  meantime  a  conductivity- 
cell  which  had  been  thoroughly  washed  the  day  before  and  in  which 
pure  alcohol  had  been  allowed  to  stand  over  night,  was  dried  with 
filtered  dry  air.  It  was  then  rinsed  several  times  with  portions  of  the 
solution  which  had  just  been  made  up,  and  finally  nearly  filled  with 
this  solution  introduced  as  shown  in  figure  23.  A  little  carbonate  is 
formed  by  opening  in  this  way  to  the  air,  but  it  is  a  very  small  quantity, 
and  in  the  course  of  a  few  days  is  entirely  precipitated  to  the  bottom 
of  the  bottle.  T  and  T'  are  filled  with  a  mixture  of  calcium  chloride 
and  soda-lime  to  protect  the  alcoholic  solution  when  the  stopcocks 
S  and  S  are  opened.  The  stoppers  in  T  and  Tf  are  of  cork  and  are 
thoroughly  paraffined.  A  system  such  as  this  remains  protected  from 
the  air  for  a  period  of  several  months. 

The  alcoholic  solution,  in  course  of  time,  becomes  colored  slightly 
yellow,  but  its  alkaline  concentration  is  apparently  not  changed,  as  can 
be  seen  by  comparing  titrations  made  against  a  standard  acid  in 
February  and  again  in  May: 

On  Feb.  25,  10  c.c.  of  standard  acid  =  8.87  c.c.  of  alkali. 

On  May   7,  10  c.c.  of  standard  acid  =  8.87  c.c.  of  alkali. 


70 


CONDUCTIVITIES   OF   ORGANIC    ACIDS 


The  bottle  containing  the  alkali  was  covered  with  a  dark  material, 
since,  in  the  presence  of  light  the  tendency  of  the  alkaline  solution  to 
become  colored  is  much  greater  than  in  the  dark. 

One  of  the  greatest  difficulties  in  connection  with  the  alcoholic  pot- 
ash method  was  that  of  temperature  changes.  The  coefficient  of  expan- 
sion of  alcohol  is  so  large  that  even  small  changes  in  the  temperature  of 
the  laboratory,  and  consequent  changes  in  temperature  of  the  solu- 
tion, will  change  quite  appreciably  the  normality  of  the  alkali. 


It  was  this  difficulty  which  led  us  to  the  use  of  an  B 

aqueous   solution  of  ammonia  with  coralline  as  j 

the   indicator,  instead    of  the    alcoholic   caustic  FIG.  24. 

potash  with  phenolphthalein  as  the  indicator. 
The  ammonia  was  prepared  by  heating  concentrated  ammonia  and 
passing  the  gas  which  was  given  off,  first  over  sticks  of  sodium  hydrox- 
ide, which  collected  a  large  part  of  the  water-vapor  and  any  carbon 
dioxide,  and  then  over  sodium,  which  absorbed  the  remainder  of  the 
water- vapor;  and  finally  into  a  weighed  quantity  of  conductivity  water 
in  a  measuring  flask  until  the  approximate  amount  of  the  gas  necessary 
to  make  a  tenth-normal  solution  was  dissolved.  This  solution  was 
titrated  against  standard  sulphuric  acid  to  obtain  its  exact  normality. 

Coralline  was  used  as  the  indicator  because  it  is  sensitive  to  the 
organic  acids,  and  is  not  sensitive  to  carbon  dioxide  except  when  the 
latter  is  present  in  fairly  large  quantity.  In  order  to  test  whether 
coralline  is  sensitive  to  small  quantities  of  carbon  dioxide,  another 
worker  in  this  laboratory  measured  out  two  equal  quantities  of  a 
standard  acid,  added  an  equal  amount  of  coralline  to  each,  and  then 
allowed  carbon  dioxide  to  bubble  through  one  of  these  solutions  for 
some  minutes.  Titrations  of  both  solutions  were  made,  and  practically 
no  effect  due  to  the  presence  of  carbon  dioxide  was  found.  Equal 
volumes  of  a  standard  acid  were  again  titrated,  this  time  after  having 


IN   ETHYL   ALCOHOL.  71 

passed  carbon  dioxide  into  one  of  the  solutions  for  a  considerable  time. 
There  was  a  small  difference  in  the  titration  values.  In  both  cases 
the  amount  of  carbon  dioxide  passed  into  the  solutions  was  infinitely 
more  than  would  ordinarily  be  present  in  such  solution  as  we  were 
titrating. 

It  was  at  first  thought  advisable  to  use  an  alcoholic  solution  of 
potassium  hydroxide  for  titration  purposes.  There  are,  however, 
several  difficulties  involved.  An  approximately  tenth-normal  solution 
of  potassium  hydroxide  in  absolute  alcohol  was  made  up  and  allowed 
to  stand  for  a  couple  of  days.  The  carbonate  settled,  leaving  a  clear 
supernatant  solution.  But  if  the  bottle  was  opened  even  for  a  very 
short  time  the  solution  became  cloudy,  and  when  poured  into  a  burette 
became  white  with  precipitated  carbonate. 

A  method  of  filtering  the  solution,  being  a  modification  of  one  pre- 
viously used  in  this  laboratory,  was  then  adopted,  together  with  an 
arrangement  for  siphoning  the  solution  out  of  the  bottle  into  the 
burette.  Figure  23  shows  the  design  of  the  filtering  apparatus.  The 
tower  T  contains  sticks  of  sodium  hydroxide  and  T'  is  partly  filled 
with  metallic  sodium.  The  former  acts  as  a  protecting  agent  to 
the  latter,  which  serves  both  for  removing  the  last  traces  of  carbon 
dioxide  and  for  drying  the  air.  B  is  a  clean  empty  bottle  which  is 
later  interchanged  with  a  bottle  filled  with  an  alcoholic  solution 
of  potassium  hydroxide  prepared  from  freshly  distilled  alcohol.  The 
tube  E  is  connected  with  suction,  so  that  dried,  purified  air  passes 
through  the  whole  system,  including  the  Gooch  funnel  F,  containing 
asbestos  previously  washed  with  an  alcoholic  solution  of  potassium 
hydroxide  and  then  pure  alcohol,  and  through  the  receiving  bottle  A. 
When  the  system  has  been  thoroughly  cleansed  with  dry  air  free  from 
carbon  dioxide,  the  stopcocks  S  are  closed,  and  the  bottle  B  is  replaced 
by  the  one  containing  alcoholic  potash.  The  stopcocks  are  then 
opened  and  suction  again  applied  to  E.  When  all  the  solution  has 
been  filtered,  A  is  removed,  and,  as  quickly  as  possible,  the  stopper 
arranged  to  connect  it  by  a  siphon  with  the  burette  (fig.  24).  It  was 
found  necessary  to  use  8  to  10  drops  of  the  solution  of  coralline  in  alco- 
hol for  each  titration.  Even  then  the  end-point  is  not  quite  as  sharp 
and  distinct  as  with  phenolphthalein. 

When  calculating  the  concentration  of  the  organic  acid  in  the  alco- 
hol from  the  values  obtained  by  titrating  against  ammonia,  it  was  found 
that  a  slightly  different  value  for  the  concentration  was  obtained  from 
that  found  from  the  titrations  against  alcoholic  caustic  potash.  We 
decided,  if  possible,  to  find  the  cause  of  this  and  to  apply  any  necessary 
corrections.  A  known  quantity  of  the  standard  sulphuric  acid  was 
titrated  against  alcoholic  potassium  hydroxide,  using  phenolphthalein  as 
the  indicator.  Several  titrations  were  made  in  every  case,  and  then  an 
equal  quantity  of  the  acid  was  titrated  against  the  base,  using  coralline. 


72  CONDUCTIVITIES   OF   ORGANIC   ACIDS 

The  results  in  the  latter  case  did  not  agree  with  those  in  the  former 
by  about  0.2  c.c.,  10  c.c.  of  acid  being  employed  in  each  case.  That 
the  difference  was  not  due  to  carbon  dioxide  which  might  have  been 
dissolved  in  the  sulphuric  acid,  can  be  seen  from  the  fact  that  the  same 
difference  appeared  in  the  titrations  with  an  organic  acid  dissolved  in 
absolute  alcohol  in  which  carbon  dioxide  is  not  very  soluble. 

It  was  found  that  if  the  same  quantity  of  phenolphthalein  or  coralline 
used  when  making  the  ordinary  titrations  was  added  either  to  pure 
alcohol  or  to  water,  and  if  these  solutions  of  the  indicators  alone  were 
titrated  against  the  alkali  and  then  back  against  the  standard  acid, 
an  appreciable  quantity  of  alkali  was  required  to  change  the  color  in 
one  direction,  and  about  as  much  of  the  standard  acid  to  change  it  in 
the  reverse  direction,  the  alkali  and  acid  being  of  very  nearly  the  same 
strength.  Corrections  for  the  amounts  of  alkali  and  acid  necessary  to 
produce  such  color  changes  were  then  applied  to  the  titration  volumes 
of  the  sulphuric  acid  and  alcoholic  potash,  when  agreement  to  within 
the  limits  of  experimental  error  was  obtained  between  the  results  for 
the  two  indicators. 

In  all  the  titrations  in  which  alcoholic  potassium  hydroxide  was  used 
its  temperature  was  recorded,  and  when  different  from  25°,  which  was 
chosen  as  the  standard  temperature,  a  volume  correction  was  applied. 
It  was  easy,  and  was  found  necessary  as  well,  to  keep  all  the  other 
solutions,  particularly  those  of  the  organic  acids  in  alcohol,  as  well  as 
the  alcoholic  potash,  at  the  standard  temperature. 

The  titration  values  of  the  ammonia  and  standard  acid  were  also 
corrected,  as  just  stated,  for  the  amounts  necessary  to  produce  color 
change,  and  the  concentration  of  the  ammonia  was  then  calculated. 
The  normality  of  1-2-4  dinitrobenzoic  acid  in  alcohol  was  determined 
from  this  standardized  ammonia,  making  the  same  corrections  as 
above;  and  it  agreed  to  within  0.2  per  cent  with  that  obtained  by 
means  of  potassium  hydroxide.  Similar  corrections  were,  therefore, 
applied  to  the  titrations  of  all  the  organic  acids. 

The  sulphuric  acid  used  to  standardize  the  alkali  was  made  up  in 
large  quantity,  and  its  normality  determined  by  the  usual  barium- 
sulphate  method. 

Owing  to  the  large  amount  of  preliminary  work  required,  it  has  been 
possible  up  to  the  present  to  make  conductivity  measurements  of  only 
9  organic  acids.  The  same  methods  of  purifying  the  acids  were 
employed  as  when  the  conductivities  of  these  acids  were  determined  in 
aqueous  solution.  In  most  cases  the  various  dilutions  were  made  up 
by  directly  weighing  the  acid. 

In  the  work  in  alcohol  it  was  necessary  to  discard  all  of  the  weaker 
organic  acids;  this,  in  spite  of  the  fact  that  our  cell  constants  were  about 
eight  times  smaller  than  those  of  Wildermann.  After  trying  acetic 
acid  several  times,  we  gave  up  hope  of  obtaining  satisfactory  results 


IN   ETHYL   ALCOHOL. 


73 


with  such  weak  acids.     Even  the  strongest  acids  with  which  we  worked 
do  not  give  a  molecular  conductivity  greater  than  unity. 

Titrations  of  the  acids  against  the  standard  alkali  were  made  simul- 
taneously with  the  conductivity  measurements  at  every  temperature. 
At  first  the  alcoholic  solution  of  the  acid  was  not  kept  at  a  constant 
temperature,  but  it  was  soon  found  that  in  order  to  obtain  comparable 
results,  and  to  avoid  the  considerable  fluctuations  of  laboratory  tem- 
perature, it  was  necessary  to  have  all  the  solutions  continuously  at  one 
temperature,  preferably  at  25°. 

RESULTS. 

In  the  following  tables  of  conductivity,  Vm  signifies  the  volume  for 
which  the  solutions  were  made  up ;  Vc  is  the  corrected  volume.  The 
corrections  applied  were  both  for  expansion  or  contraction  of  the  alco- 
hol, and  for  change  in  the  concentration  of  the  acid  due  to  formation 
of  ester.  Molecular  conductivity  was  calculated  in  the  usual  manner. 
Temperature  coefficients  and  percentage  temperature  coefficients  are 
expressed  for  10  degrees.  The  specific  conductivities  of  the  alcohol  as 
given  are  all  multiplied  by  103.  They  are  really  of  the  magnitude  10~7. 

TABLE  23. — Molecular  conductivities,  temperature  coefficients,  etc.,  of  certain  acids. 


Malonic  acid. 

O'Chlorobenzoic  acid. 

vm 

Vc 

Mr 

Time  of  reading. 

vm 

F. 

Ht> 

Time  of  reading. 

15° 

16° 

8 

8.12 

0.0190 

Apr.  16,    lh  00m  p.m. 

8 

8.14 

0.01303 

Mar.  18,  10h  40™  p.m. 

32 

32.9 

0.0434 

16,    1 

05    p.m. 

32 

33.1 

0.01530 

18,  10 

45    p.m. 

128 

129.3 

0.0775 

17,12 

20    p.m. 

128 

129.5 

0.0279 

19,  10 

50   p.m. 

512           512.8 

0.2533 

17,12 

45    p.m. 

512 

513.8 

0.1330 

19,  11 

15    p.m. 

Alcohol      Sp.  cond. 

0.000254 

16,    1 

10    p.m. 

Alcohol 

Sp.  cond. 

0.000531 

18,11 

20    p.m. 

Alcohol      Sp.  cond. 

0.000246 

17,12 

45    p.m. 

Alcohol 

Sp.  cond.     0  .  000540 

19,11 

30   p.m. 

25° 

9P 

8                8.13 

0.0237 

16,    2 

35    p.m. 

8 

8.16       0.0159 

18,11 

20    a.m. 

32              33.20 

0.0555 

16,    2 

40    p.m. 

32 

33.60 

0.0198 

18,11 

45    a.m. 

128            129.5 

0.0985 

17,    2 

40    p.m. 

128 

129.7 

0.0371 

19,12 

35    p.m. 

512 

514.9 

0.3160 

17,    2 

45    p.m. 

512 

516.5 

0.1714 

19,12 

40    p.m. 

Alcohol 

Sp.  cond. 

0.000257 

16,    2 

45    p.m. 

Alcohol 

Sp.  cond. 

0.000578 

18,11 

50    a.m. 

Alcohol 

Sp.  cond. 

0.000249 

17,    2 

50    p.m. 

Alcohol 

Sp.  cond. 

0.000622 

19,12 

45    p.m. 

35° 

35° 

8 

8.18 

0.0319 

16,    4 

30    p.m. 

8 

8.28 

0.0197 

18,    4 

45    p.m. 

32 

33.7 

0.0737 

16,    4 

35    p.m. 

32 

34.9 

0.0271 

18,    4 

50    p.m. 

128 

129.6 

0.1351 

17,    4 

20    p.m. 

128 

129.8 

0.0555 

19,    4 

55    p.m. 

512 

518.1 

0.4338 

17,    4 

25    p.m. 

512 

519.9 

0.2497 

19,    5 

00    p.m. 

Alcohol 

Sp.  cond.     0.000258 

16,    4 

40    p.m. 

Alcohol 

Sp.  cond. 

0.000637 

18,    5 

00    p.m. 

Alcohol 

Sp.  cond. 

0.000246 

17,    4 

30    p.m. 

Alcohol 

Sp.  cond. 

0.000711 

19,    5 

10    p.m. 

Temperature  coefficients. 

Temperature  coefficients. 

vm 

15°  to  25° 

25°  to  35° 

vm 

15°  to  25° 

25°  to  35° 

Con.  unit.       P.  ct. 

Con.  unit. 

P.  ct. 

Con.  unit. 

P.  ct. 

Con.  unit. 

P.  ct. 

8 

0.0046 

24.5 

0.0075 

32.2 

8 

0.00281 

21.1 

0.0040 

25.9 

32 

0.0113 

26.7 

0.0165 

30.7 

32 

0.00407 

27.5 

0.0060 

31.2 

128 

0.0207 

26.9 

0.0360 

37.0 

128 

0.00903 

32.8 

0.0181 

49.2 

512 

0.0613         24.2 

0.1145 

36.4 

512 

0.0374 

28.1 

0.0759 

44.6 

74 


CONDUCTIVITIES   OF   ORGANIC    ACIDS 


TABLE  23. — Molecular  conductivities,  temperature  coefficients,  etc.,  of  certain  acids — Continued. 


p-Chlorobenzoic  acid. 

o-Nitrobenzoic  acid. 

vm          vc 

MP 

Time  of  reading. 

vm 

Vc 

Me 

Time  of  reading. 

15° 

15° 

8 

10.0 

0.0017 

Mar.26,12h30mp.m. 

8 

8.21 

0.00785 

Apr.  23,  12h  35m  p.m. 

32 

33.69 

0.0082 

26,12 

35    p.m. 

32 

33.19 

0.0204 

23,12 

40    p.m. 

128 

129.7 

0.0157 

Apr.    1,  12 

30    p.m. 

128 

129.3 

0.0460 

27,12 

30    p.m. 

512 

514.8 

0.1263 

1,  12 

35    p.m. 

512 

512.5 

0.1788 

27,  12 

25    p.m. 

Alcohol 

Sp.  cond. 

0.000585 

Mar.  26,  12 

40    p.m. 

Alcohol 

Sp.  cond. 

0.000232            23  12 

50    p.m. 

Alcohol 

Sp.  cond. 

0.000586 

Apr.    1,12 

45    p.m. 

Alcohol 

Sp.  cond. 

0.000227            27  12 

35    p.m. 

25° 

%o 

8 

10.08 

0.0025 

Mar.  26,    2 

50    p.m. 

8 

8.27 

0.00937 

23,    2 

20    p.m. 

32 

34.18 

0.0117 

26,    2 

55    p.m. 

32 

34.30 

0.0253 

23,    2 

40    p.m. 

128 

130.3 

0.0189 

Apr.    1,    3 

10    p.m. 

128 

129.5 

0.0477 

27,    2 

40    p.m. 

512 

520.0 

0.1547 

1,    3 

15    p.m. 

512 

517.0 

0.2452 

27,    2 

50    p.m. 

Alcohol 

Sp.  cond. 

0.000650  Mar.  26,    3 

00    p.m. 

Alcohol 

Sp.  cond. 

0.000242            23,    2 

50    p.m. 

Alcohol 

Sp.  cond. 

0.000656  Apr.    1,    3 

20   p.m. 

Alcohol 

Sp.  cond. 

0.000238            27,    2 

50    p.m. 

35° 

35° 

8 

10.11 

0.0035 

Mar.  26,    4 

40    p.m. 

8 

8.27 

0.0120                23,    4 

30    p.m. 

32 

34.9 

0.0160 

26,    4 

45    p.m. 

32 

34.95 

0.0337                 23, 

35    p.m. 

128 

130.9 

0.0270 

Apr.    3,    4 

10    p.m. 

128 

129.6 

0.0734                27, 

30    p.m. 

512 

522.7 

0.1853 

3,    4 

15    p.m. 

512 

518.5 

0.2877                27, 

35    p.m. 

Alcohol 

Sp.  cond. 

0.000753 

Mar.  26,    4 

50    p.m. 

Alcohol 

Sp.  cond. 

0.000237            23, 

45    p.m. 

Alcohol 

Sp.  cond. 

0.000827 

Apr.    3,   4 

20    p.m. 

Alcohol 

Sp.  cond. 

0.000232            27, 

40    p.m. 

Temperature  coefficients. 

Temperature  coefficients. 

vm 

15°  to  25° 

25°  to  35° 

Vm 

15°  to  25° 

25°  to  35° 

Con.  unit.           P.  ct. 

Con.  unit. 

P.  ct. 

Con.  unit. 

P.  ct. 

Con.  unit. 

P.  ct. 

8 

0.00052           42.2 

0.00079 

39.8 

8 

0.00142 

18.56 

0.0025 

27.45 

32 

0.00315           40.2 

0.00382 

34.7 

32 

0.0040 

20.31         0.0073 

30.90 

128 

0.00314           19.9 

0.0079 

42.4 

512 

0.0642 

35.94         0.0413 

17.00 

512 

0  0268             27  2 

0  0293 

1Q    9 

iy  .  _ 

p-Nitrobenzoic  acid. 

p-Bromobenzoic  acid. 

vm 

Vc 

Mr 

Time  of  reading. 

vm 

Vc 

MP 

Time  of  reading. 

15° 

15° 

8 

8.147 

0.00264 

Apr.  28,  Ilh40ma.m. 

32 

32.96 

0.0102 

Apr.  21,  12h05mp.m. 

32 

32.57 

0.01252 

28,11 

45    a.m. 

128 

129.2 

0.0516 

21,12 

10    p.m. 

128 

129.1 

0.0349 

29,11 

10    a.m. 

512 

512.8 

0.1417 

22,  12 

05    p.m. 

512 

512.8 

0.1651 

29,  11 

15    a.m. 

Alcohol 

Sp.  cond. 

0.000237 

21,12 

15    p.m. 

Alcohol 

Sp.  cond. 

0.000217 

28,11 

50    a.m. 

Alcohol 

Sp.  cond. 

0.000236 

22,12 

15    p.m. 

Alcohol 

Sp.  cond. 

0.000214 

29,11 

20    a.m. 

Alcohol 

Sp.  cond. 

0.000231 

22,12 

20    p.m. 

25° 

25° 

8 

8.24 

0.00353 

28,    2 

00    p.m. 

32 

33.61 

0.0151 

21,    2 

45    p.m. 

32 

33.3 

0.0147 

28,    2 

10    p.m. 

128 

129.3 

0.0570 

21,    2 

50    p.m. 

128 

129.2 

0.0418 

29,12 

25    p.m. 

512 

517.6 

0.1814 

22,    2 

35    p.m. 

512 

Alcohol 
Alcohol 

517.5 
Sp.  cond. 
Sp.  cond. 

0.1976 
0.000216 
0.000262 

29,  12 
28,    2 
29,12 

30    p.m. 
20    p.m. 
35    p.m. 

Alcohol 
Alcohol 
Alcohol 

Sp.  cond. 
Sp.  cond. 
Sp.  cond. 

0.000237 
0.000227 
0.000227 

21,    2 
22,    2 
22,    2 

50    p.m. 
45    p.m. 
50    p.m. 

35° 

35° 

8 

8.27 

0.0047 

28, 

45    p.m. 

32 

34.56 

0.0214 

21,    4 

50    p.m. 

32 

34.33 

0.0200 

29, 

50    p.m. 

128 

129.5 

0.0785 

21,    4 

55    p.m. 

128 

129.5 

0.0559 

28, 

00    p.m. 

512 

520.5 

0.2399 

22,    4 

35    p.m. 

512 

Alcohol 
Alcohol 

518.9 
Sp.  cond. 
Sp.  cond. 

0.2637 
0.000264 
0.000233 

29, 

28, 
29, 

05    p.m. 
55    p.m. 
15    p.m. 

Alcohol 
Alcohol 
Alcohol 

Sp.  cond. 
Sp.  cond. 
Sp.  cond. 

0.000232 
0.000219 
0.000216 

21,    4 
22,    4 
22,    4 

00    p.m. 
40    p.m. 
45    p.m. 

Temperature  coefficients. 

Temperature  coefficients. 

v_ 

i  ^°  +/\  *>^° 

OCO   4.-    OKO 

'  n 

1O     tO  ^5O 

Zo    to  00 

TT 

1  £°  4-^   O£° 

OKO   A.-.    OCO 

c 

P.  ct. 

P*t 

"  m 

lo    to  -•> 

/o   to  oo 

8 

o.ooos' 

30.7 

C071.  117111. 

0.0012 

.  ct. 
34.9 

Con.  unit. 

P.  ct. 

Con.  unit. 

P.  ct. 

32 

0.0019 

15.4 

0.0046 

32.3 

32 

0.0045 

45.5 

0.0055 

38.2 

128 

0.0068 

19.6 

0.0138 

33.3 

128 

0.0053 

10.3 

0.0212 

37.5 

512 

0.0307 

18.6 

0.0607 

29.9 

512 

0.0380 

26.8 

0.0565 

31.4 

IN   ETHYL   ALCOHOL. 


75 


TABLE  23. — Molecular  conductivities,  temperature  coefficients,  etc.,  of  certain  acids — Continued. 


1,2,4  Dinitrobenzoic  acid. 

1 

,  2,  4  Dihydroxybenzoic  acid  —  Continued. 
Temperature  coefficients. 

vm 

Vc 

Me 

Time  of  reading. 

8 
32 
128 
Alcohol 
Alcohol 

8 
32 
128 
Alcohol 
Alcohol 

8 
32 
128 
Alcohol 
Alcohol 

8.13 
33.62 
133.50 
Sp.  cond. 
Sp.  cond. 

8.24 
33.62 
133.50 
Sp.  cond. 
Sp.  cond. 

8.24 
33.62 
133.5 
Sp.  cond. 
Sp.  cond. 

16° 

0.0379 
0.0964 
0.2556 
0.000882 
0.000936 
25° 
0.0481 
0.0848 
0.1670 
0.000991 
0.000935 
35° 
0.05879 
0.10512 
0.20043 
0.001133 
0.00123 

Feb.  27,  12h  00™  m. 
Mar.    3,  12h10mp.m. 
Feb.  27,  12  25    p.m. 
27,12  25    p.m. 
Mar.    3,  10  35    a.m. 

Feb.  27,  12  40    p.m. 
27,    2   15    p.m. 
Mar.    3,  12   30    p.m. 
Feb.  27,    2  25    p.m. 
Mar.    3,  10  40    a.m. 

Feb.  27,    4  45    p.m. 
27,   3  45    p.m. 
Mar.    3,    3  45    p.m. 
Feb.  27,    3  50    p.m. 
Mar.    3,    3  45    p.m. 

Vm 

15°  to  25° 

25°  to  35° 

8 
32 
128 
512 

Con.  unit. 
0.0014 
0.004 
0.010 
0.0445 

P.ct. 

21.9 
26.5 
59.1 
44.3 

Con.  unit. 
0.0023 
0.0054 
0.0117 
0.0539 

P.ct 
28.3 
28.3 
43.4 
37.2 

Tetrachlorphthalic  acid. 

Vm 

y 

M» 

Time  of  reading. 

16 
64 
256 
1024 
Alcohol 
Alcohol 

16 
64 
256 
1024 
Alcohol 
Alcohol 

16 
64 
256 
1024 
Alcohol 
Alcohol 

16.08 
64.06 
258.9 
1027.0 
Sp.  cond. 
Sp.  cond. 

16.21 
64.06 
259.3 
1036.0 
Sp.  cond. 
Sp.  cond. 

16.21 

64.06 
260.0 
1043.0 
Sp.  cond. 
Sp.  cond. 

15° 
0.0543 
0.1011 
0.1294 
0.3208 
0.000543 
0.000554 
25° 
0.0639 
0.1198 
0.1541 
0.3860 
0.000616 
0.000637 
35° 
0.0770 
0.1461 
0.1893 
0.4960 
0.000711 
0.000742 

Mar.  20,  Ilh50ma.m. 
20,  12  00  m. 
21,12   15    p.m. 
21,12  25    p.m. 
20,12  05    p.m. 
21,12  45    p.m. 

20,  12  50    p.m. 
20,    1   00   p.m. 
21,    2  00    p.m. 
21,    2   10    p.m. 
20,12  40    p.m. 
21,    2   15    p.m. 

20,    4  30    p.m. 
20,    4  30    p.m. 
21,    4  00    p.m. 
21,    4   10    p.m. 
20,    3  25    p.m. 
21,    4   10    p.m. 

1,  2,  4  Dihydroxybenzoic  add. 

Vm 

Vc 

Mi 

Time  of  reading. 

8 
32 
128 
512 

Alcohol 
Alcohol 

8 
32 
128 
512 
Alcohol 
Alcohol 

8 
32 
128 
512 
Alcohol 
Alcohol 

9.99 
33.0 
129.1 
514.9 
Sp.  cond. 
Sp.  cond. 

10.06 
33.01 
129.3 
517.5 
Sp.  cond. 
Sp.  cond. 

10.10 
33.01 
129.6 
520.2 
Sp.  cond. 
Sp.  cond. 

15° 
0.0080 
0.0155 
0.0171 
0  .  1008 
0.000551 
0.000613 
25° 
0.0098 
0.0197 
0.0272 
0.1464 
0.000631 
0.000682 
35° 
0.0126 
0.0254 
0.0391 
0.2018 
0.000735 
0.000791 

Mar.  24,  llh  00™  a.m. 
24,  11   05    a.m. 
25,  11   45    a.m. 
25,11   50    a.m. 
24,11    15    a.m. 
25,11   55    a.m. 

24,    2  30    p.m. 
24,    2  35    p.m. 
25,    2  35    p.m. 
25,    2  40    p.m. 
24,    2  40    p.m. 
25,   2  50   p.m. 

24,    4  20    p.m. 
24,    4  25    p.m. 
25,    4   15    p.m. 
25,    4  20    p.m. 
24,    4  35    p.m. 
25,    4   25    p.m. 

Temperature  coefficients. 

Vm 

15°  to  25° 

25°  to  35° 

16 
64 
256 
1024 

Con.  unit. 
0.0091 
0.0187 
0.0243 
0.0615 

P.  ct. 
16.9 
18.5 
18.1 
19.3 

Con.  unit. 
0.0138 
0.0262 
0.0342 
0.1057 

P.  ct. 

18.9 
21.9 
22.5 
27.6 

TABLE  24. — Changes  in  concentration. 


Malonic  acid. 

Normality. 

Time. 

Observed. 

Calcu- 
lated. 

De- 
crease. 

P.ct. 

Apr.  16,  12h  45m  a.m. 

0.1232 

0.1250 

1.44 

16,    2  30    p.m. 

0.1229 

" 

1.68 

16,    4  00    p.m. 

0.1222 

" 

2.24 

16,12  50    p.m. 

0.03039 

0.03125 

2.75 

16,    2  30    p.m. 

0.03008 

" 

3.75 

16,    4   10    p.m. 

0.02966 

" 

5.09 

17,12  00    m. 

0.00773 

0.007812 

1.05 

17,    2  30    p.m. 

0.00772 

" 

1.18 

17,   4  00    p.m. 

0.00771 

" 

1.31 

17,  12   10    p.m. 

0.001950 

0.001953 

0.16 

17,    2  40    p.m. 

0.001942 

" 

0.57 

17,    4   10    p.m. 

0.001930 

1.18 

76 


CONDUCTIVITIES   OF   ORGANIC   ACIDS 


TABLE  24. — Changes  in  concentration — Continued. 


o-Chlorobenzoic  acid. 

p-Ntirobenzoic  acid. 

Time. 

Normality. 

Time. 

Normality. 

Observed. 

Calc. 

De- 
crease. 

Observed. 

Calc. 

De- 
crease. 

Mar.  18,  l^OC^a.m. 
18,12   15    p.m. 
18,    4   15    p.m. 
18,    4  40    p.m. 
18,11  00    a.m. 
18,11  50    a.m. 
18,    3  30    p.m. 
18,    4  30    p.m. 
19,11  45    a.m. 
19,12  30    p.m. 
19,    2  30    p.m. 
19,   4  00    p.m. 
19,  11  35    a.m. 
19,12  30    p.m. 
19,    2  50    p.m. 
19,    4  45    p.m. 

0.1228 
0.1225 
0.1210 
0.1208 
0.03018 
0.02976 
0.02914 
0.02862 
0.00772 
0.00771 
0.00770 
0.00769 
0.001946 
0.001936 
0.001932 
0.001923 

0.1250 
0.03125 
0.007812 
0.001953 

P.ct. 
1.76 
2.00 
3.20 
3.36 
3.43 
4.77 
6.75 
8.42 
1.18 
1.30 
1.43 
1.56 
0.36 
0.87 
1.08 
1.54 

Apr.  21,  12hOOmm. 
21,    2  30    p.m. 
21,    4  45    p.m. 
22,12  00    m. 
22,    2  30    p.m. 
22,    4  45    p.m. 
22,12  00    m. 
22,    2  30    p.m. 
22,    4  30    p.m. 

0.03034 
0.02976 
0.02893 
0.007738 
0.007734 
0.007721 
0.001950 
0.001932 
0.001921 

0.03125 
0.007812 
0.001953 

P.  ci. 
2.91 
.77 
.43 
.95 
.00 
.17 
.16 
.08 
.64 

1,2,4  Dinitrobenzoic  acid. 

Feb.  27,  12h30mp.m. 
27,    2  30    p.m. 
27,    4  30    p.m. 
27,12  20    p.m. 
27,    4  00    p.m. 
27,    5  00    p.m. 
Mar.    3,  12  35    p.m. 
3,    2  30    p.m. 
»4,  10  00    a.m. 

0.1230 
0.1213 

0.02974 

0.00749 
0.00764 

0.1250 
0.03125 
0.007812 

1.60 
2.96 
2.96 
4.83 

4.13 
4.13 
2.20 

p-Chlorobenzoic  add. 

Mar.  26,  12h  30™  p.m. 
26,    2  00    p.m. 
26,    4   15    p.m. 
26,12  40    p.m. 
26,    2  00    p.m. 
26,    4  30    p.m. 
Apr.     1,  12  30    p.m. 
1,    2  20    p.m. 
3,    3  30    p.m. 
3,    4   10    p.m. 
3,    5  00    p.m. 
1,12  35    p.m. 
1,    3   15    p.m. 
3,  12  30    p.m. 
3,    3  30    p.m. 
3,    4   15    p.m. 

0.0996 
0.09919 
0.0989 
0.0297 
0.0292 
0.0286 
0.00771 
0.007691 
0.00767 
0.00766 
0.00763 
0.001942 
0.001923 
0.001916 
0.001913 
0.001903 

0.1250 
0.03125 
0.007812 

0.001953 

20.32 
20.65 
20.90 
4.96 
6.55 
8.42 
.31 
.55 
.82 
.95 
.33 
.57 
.54 

2.05 
2.57 

•This  titration  was  made  with  the  solution  after  it 
had  stood  in  the  cell  over  night. 

1,2,4  Dihydroxybenzoic  acid. 

Mar.24,  Ilh15m  a.m. 
24,12  30    a.m. 
24,    2  30    a.m. 
24,    4  30    a.m. 
24,  12  25    p.m. 
24,  12  45    p.m. 
24,    2  30    p.m. 
24,    4  30    p.m. 
25,12  00    m. 
25,    2  30    p.m. 
25,    4   15    p.m. 
25,  12   10    p.m. 
25,    2  40    p.m. 
25,    4  30   p.m. 

0.10008 
0.09940 
0.09899 
0.09873 
0.03034 
0.03029 
0.03029 
0.03013 
0.00774 
0.00773 
0.00771 
0.00194 
0.00193 
0.00192 

0.1250 
0.03125 

0.007812 
0.001953 

19.94 
20.50 
20.82 
21.02 
2.91 
3.08 
3.08 
3.59 
0.92 
1.03 
1.31 
>   0.67 
1.18 
1.69 

p-Bromobenzoic  acid. 

Apr.  28,  Ilh30ma.m. 
28,    2  45    p.m. 
28,    4  40    p.m. 
28,11   40    a.m. 
28,    2  50    p.m. 
28,    4  50    p.m. 
29,  11   00    a.m. 
29,  12   15    p.m. 
29,    4  00    p.m. 
29,11    10    a.m. 
29,12  25    p.m. 
29,   4   10    p.m. 

0.12275 
0.12129 
0.12088 
0.03070 
0.03007 
0.02914 
0.00774 
0.00773 
0.00771 
0.001950 
0.001932 
0.001927 

0.1250 
0.03125 
0.00781 
0.001953 

1.80 
2.97 
3.30 
1.76 
3.78 
6.76 
0.89 
1.02 
1.28 
0.16 
1.08 
1.34 

Tetrachlorphthalic  acid. 

Mar.  20,  llh  50™  a.m. 
20,12   50    p.m. 
20,    3  20    p.m. 
20,    4  30    p.m. 
20,12  00    m. 
20     1   00    p.m. 
20,    3  30    p.m. 
20,    4  30    p.m. 
21,  12   15    p.m. 
21,    2  00    p.m. 
21,    4  00    p.m. 
21  12  25    p.m. 
21,    2   10    p.m. 
21,    4   10    p.m. 

0.06218 
0.06168 

0.01561 

0.00386 
0.00385 
0.00384 
0.000973 
0.000965 
0.000958 

0.0625 
0.01566 

0.00391 
0.000976 

0.52 
1.32 

0.32 

1.28 
1.53 
.79 
0.31 
.13 
.85 

o-Nitrobenzoic  acid. 

Mar.  23,  12h  15""  p.m. 
23,  12  30    p.m. 
23,    4  30    p.m. 
23,  12  30    p.m. 
23,    2  30    p.m. 
23,    4  30    p.m. 
27,  12  00    m. 
27,    2  30    p.m. 
27,    4  00    p.m. 
27,12   10    p.m. 
27,    2  40    p.m. 
27,    4   10    p.m. 

0.12176 
0.12088 
0.12088 
0.03013 
0.02914 
0.0286 
0.00773 
0.00772 
0.00771 
0.001951 
0.001934 
0.001928 

0.1250 
0.03125 
0.00781 
0.001953 

2.60 
3.30 
3.30 
3.52 
6.75 
8.48 
1.03 
1.17 
1.29 
0.10 
0.98 
1.28 

IN   ETHYL   ALCOHOL.  77 

DISCUSSION  OF  RESULTS. 

It  will  be  noted  in  tables  23  and  24  that  1,  2,  4  dinitrobenzoic  acid 
shows  irregularity  in  titration  values.  The  conductivity  of  this  acid 
was  determined  before  we  began  to  keep  the  solutions  used  in  titrating 
at  a  constant  temperature.  With  all  other  acids  the  results  show  that 
with  increase  in  time  a  greater  amount  of  esterification  has  taken  place; 
that  is,  the  normality  of  the  acid  has  become  less.  The  amount  of 
ester  formed  in  a  given  time  depends  upon  the  nature  of  the  acid. 

Since  each  dilution  was  made  up  independently  of  the  others,  that 
is,  by  direct  weight,  it  is  interesting  to  note  that  the  proportion  of 
ester  formed  in  the  less  dilute  solutions  is  much  greater  than  in  the  more 
dilute.  Indeed,  in  some  cases  there  is  practically  no  ester  formed  in 
the  y^-g-  and  -5^  solutions. 

As  has  already  been  stated,  none  of  the  conductivities  is  greater  than 
unity;  and,  consequently,  the  molecular  conductivity  of  the  alcohol  for 
each  dilution  is  relatively  quite  large,  the  correction  for  this  factor  being 
in  some  cases  as  much  as  70  per  cent  of  the  total  conductivity.  It  can 
be  seen  from  the  tables  that  the  conductivity  of  the  alcohol  alone  varies 
considerably,  usually  increasing  appreciably  with  time.  Some  of  the 
conductivities  of  the  alcohol  increase  with  rise  in  temperature,  some 
actually  decrease,  while  others  remain  very  nearly  constant.  We  can 
offer  no  explanation  for  this  lack  of  uniform  variation,  except  to  call 
attention  to  the  several  factors  which  might  affect  the  conductivity 
of  the  pure  solvent.  One  might  be  the  absorption  by  the  alcohol  of 
traces  of  various  gases  or  water- vapor  from  the  atmosphere.  This, 
however,  ought  to  be  a  negligible  factor,  since  our  cells  were  very  nearly 
filled  and  were  tightly  closed  with  ground-glass  stoppers. 

The  decomposition  effects  brought  about  by  the  platinum  electrodes 
may  be  an  important  factor.  (Compare  here  the  work  of  Wildermann 
and  others  on  this  question.)  It  is  evident  that  the  electrodes  do  have 
some  effect,  since  fresh  alcohol  just  taken  from  the  bottle  does  have  a 
fairly  uniform  conductivity.  Part  of  the  effect  with  alcohol  which 
stood  in  the  cell  over  night  might  be  due  to  the  solubility  of  the  glass 
cells.  This,  however,  is  not  at  all  probable,  since  our  cells  have  been 
in  constant  use  in  this  laboratory  for  three  years,  and  hard  glass  is  not 
very  soluble  in  alcohol. 

The  conductivities  of  some  of  the  solutions,  and  curiously  enough  of 
the  more  dilute  solutions,  vary  to  a  much  smaller  extent,  with  time, 
than  does  the  conductivity  of  the  pure  alcohol. 

It  will  be  recalled  that  Wakemann  plotted  curves  of  conductivity  of 
the  organic  acids  against  percentage  alcohol  (see  fig.  21),  and  on  extend- 
ing the  curves  in  the  direction  of  100  per  cent  alcohol  they  apparently 
approached  zero  conductivity  as  a  limit.  As  can  be  seen  from  our 
results,  the  conductivities  do  not  actually  approach  zero,  but  a  number 
less,  and  usually  very  much  less  than  unity. 


78 


CONDUCTIVITIES   OF   ORGANIC   ACIDS 


One  of  the  most  interesting  facts  developed  in  this  work  is  the  very 
large  percentage  temperature  coefficients  of  conductivity  of  the  organic 
acids  in  alcohol.  These  range  for  10°  from  15  to  50  per  cent. 

There  is  often  a  rapid  increase  in  the  conductivity  of  the  organic  acids 
with  increase  in  dilution,  yet  certain  of  the  acids  behave  in  just  the 
opposite  manner — e.  g.,  o-chlorbenzoic  acid  and  p-nitrobenzoic  acid. 

Our  results  seem  to  suggest  the  following  possibilities,  if  we  take 
into  account  the  work  done  here  on  the  organic  acids  in  aqueous  solu- 
tions; that  there  is  much  greater  alcoholation  than  hydration,  and  this 
is  decreased  with  rise  in  temperature.  The  work  already  done  in  this 
laboratory  renders  this  highly  improbable.  The  alcoholates  may  be 

more  unstable  with  rise  in  temperature 
than  the  hydrates,  but  water  seems  to 
have  in  general  far  more  power  to  com- 
bine with  dissolved  substances  than 
alcohol. 


0.16 
0.15 
0.14 
0.13 
x.0.12 
1  OLII 
'H  o.io 

T5 

§  0.09 

0 

u  0.08 
1  -0-07 
I  0.06 
0.05 
0.04 
0.03 
0.02 
0.01 
0.0 

/ 

I 

// 

// 

/// 

1 

1 

II 

' 

I/ 

II 

I/ 

I/ 

1 

j 

^ 

j 

^ 

^^ 

3                        1                        Z                       3 

Log  volume 
FIG.  25. — Malonic  acid. 


Log 
FIG.  26. — p-Chlorobenzoic  acid. 


If  dissociation  in  alcoholic  solutions  increases  with  rise  in  tempera- 
ture, it  might  account  for  the  large  temperature  coefficients  of  conduc- 
tivity in  such  solutions ;  but  this  again  seems  highly  improbable.  The 
greater  expansion  of  the  alcohol  with  rise  in  temperature  would  allow  a 
freer  movement  of  the  ions,  and  this  doubtless  is  of  some  significance. 

A  method  for  determining  the  dissociation  of  the  organic  acids  in 
alcohol  (somewhat  similar  to  that  used  with  aqueous  solutions)  will, 
it  is  hoped,  be  worked  out  in  the  investigation  of  this  subject  which  is 
to  follow  this  preliminary  one.  It  will  involve  the  study  in  alcohol  of 


IN   ETHYL   ALCOHOL. 


79 


0.25 
0.20 
0.15 


the  conductivity  of  some  salts  of  the  acids,  as  well  as  of  hydrochloric 
acid  and  the  chlorides  corresponding  to  these  salts. 

The  increase  in  conductivity  with  increase  in  volume  is  shown 
graphically  in  figures  25  and  26.  The  increase  in  conductivity  with 
rise  in  temperature  can  be  seen  from  figs.  27  and  28.  In  the  latter  case 
the  curves  have  very  much  the  appearance  of  those  in  aqueous  solu- 
tions. This  suggests  that  perhaps  the  increase  in  molecular  conduc- 
tivity in  alcohol  with  rise  in  temperature  is  a  parabolic  function,  as 
in  aqueous  solutions,  and  that  the  Euler  equation,  nv  =  fj,Q+at  —  bP, 
applies  to  both. 


x 


0.15 


15  25" 

Temperature 

FIG.  27. — Malonic  acid. 


5  15  25 

Temperature 

FIG.  28. — p-Chlorobenzoic  acid. 


This  will  be  tested  in  the  later  work  by  determining  the  conduc- 
tivities of  some  of  the  acids  at  temperatures  other  than  the  three 
already  named,  and  comparing  the  results  obtained.  The  most  striking 
feature  of  the  conductivities  of  the  same  acids  in  water  is  their  very 
small  value.  When  we  consider  the  relative  powers  of  alcohol  and 
water  to  dissociate  salts,  the  above  fact  does  not  at  present  seem  to 
admit  of  any  very  satisfactory  explanation.  Alcohol  has  from  one- 
fourth  to  one-fifth  the  dissociating  power  of  water,  as  shown  by  their 
dissociation  of  salts.  With  the  organic  acids  the  conductivities  in 
alcohol  are  often  several  hundred  times  smaller  than  in  water.  It  is 
hoped  that  the  further  work  which  is  now  in  progress  in  this  laboratory 
on  this  problem  may  throw  some  light  on  this  relation. 


CHAPTER  IV. 


THE  CONDUCTIVITY  AND  VISCOSITY  OF  SOLUTIONS  OF  POTASSIUM 

IODIDE  AND  SODIUM  IODIDE  IN  MIXTURES  OF  ETHYL 

ALCOHOL  AND  WATER. 

BY  E.  P.  WIGHTMAN,  P.  B.  DAVIS,  AND  A.  HOLMES. 

A  brief  review  of  the  conductivity  and  viscosity  work  in  non-aqueous 
and  mixed  solvents,  during  the  past  twelve  years,  is  contained  in  the 
last  chapter  of  this  monograph.1  All  discussion  of  this  work  can, 
therefore,  be  omitted  here. 

EXPERIMENTAL. 
PURE  ANHYDROUS  ALCOHOL. 

Pure  anhydrous  alcohol  was  obtained  in  the  following  manner :  The 
ordinaiy  95  per  cent  ethyl  alcohol  was  heated  for  three  days  with  lime 
in  a  copper  vessel  connected  with 
a  reflux  condenser.  A  cooling  coil 
in  the  neck  of  the  vessel  brought 
about  rapid  condensation,  thus 
acting  as  a  safety  device,  so  that 
there  was  no  danger  in  keeping 
the  alcohol  constantly  heated 
during  the  day  without  close  at- 
tention. 

In  distilling  the  alcohol,  a  block- 
tin  condenser  connected  with  the 
copper  vessel  by  means  of  a 
ground-brass  joint  (see  fig.  29)  was 
used.  In  this  way  the  ordinary 
cork  stopper  was  avoided,  and 
the  alcohol  vapor  came  in  contact 
only  with  a  metal  surface  before  being  condensed.  The  distillate  was 
received  into  large  (glass-stoppered)  Jena  glass  bottles. 

Specific-gravity  determinations  showed  this  to  contain  from  0.1  per 
cent  to  0.07  per  cent  of  water.  It  was,  therefore,  heated  a  second  time 
for  three  or  four  days  with  fresh  lime  and  then  redistilled.  The  dis- 
tillate obtained  in  this  manner  had  a  specific  gravity  from  about  0.78511 
to  0.78516,  usually  nearer  the  former  value,  which  corresponds  to  a 
percentage  of  99.98  per  cent  alcohol. 

SPECIFIC-GRAVITY  DETERMINATIONS. 

Special  care  was  taken  in  the  determination  of  densities.  Two 
pycnometers  (fig.  30),  very  nearly  alike,  were  used  in  the  case  of 


XJ 

FIGS.  29  and  30. 


30 


80 


iSee  also  Carnegie  Inst,  Wash.  Pubs.  Nos.  80  and  180. 


SOLUTIONS    OF    SALTS   IN   ETHYL   ALCOHOL    AND    WATER.      81 

pure  alcohol.  They  were  similar  in  shape  to  those  employed  in  earlier 
work,  but  were  nearly  twice  as  large,  having  a  volume  capacity 
somewhat  over  20  c.c.,  and  the  capillary  was  of  0.5  mm.  bore.  By 
using  one  of  these  as  a  tare  against  the  other,  effects  caused  by 
changes  in  atmospheric  conditions  were  avoided.  It  may  be  said  here 
also  that  in  all  weighings  the  load  was  weighed  on  each  end  of  the 
balance  beam,  and  that  the  final  weight  represented  the  mean  value  of 
the  two.  For  all  other  specific-gravity  determinations  smaller  pyc- 
nometers  with  capacities  of  about  10  c.c.,  were  employed,  and  were 
weighed  directly,  as  in  the  previous  case,  on  each  end  of  the  beam. 

Corrections  were  always  applied  to  the  apparent  weights  of  the  con- 
tents of  the  pycnometers  in  order  to  reduce  them  to  the  vacuum 
standard.  For  this  purpose  a  record  was  kept  of  the  height  of  the 
barometer  and  the  temperature  of  the  balance-room  at  the  time  of 
weighing.  The  buoyancy  correction  was  afterwards  determined  by 
means  of  table  22,  page  37  of  Circular  No.  19  of  the  Bureau  of  Standards. 
The  capacities  of  the  pycnometers  were  found  in  the  usual  manner, 
with  the  addition  of  the  corrections  just  mentioned,  at  15°,  25°,  and 
35°  C.  Moreover,  the  pycnometers  were  reset  and  re  weighed  twice  at 
each  temperature,  in  order  to  be  sure  that  the  capacities  were  correct. 

MIXED  SOLVENTS. 

The  mixed  solvents  were  made  up  in  percentages  by  weight  of  alcohol 
and  water.  These  percentages  were  found  from  the  density  tables  on 
pages  6  and  7  of  Circular  No.  19  of  the  Bureau  of  Standards.  The 
making  of  the  mixtures  of  alcohol  and  water  on  a  weight  basis  was  by  a 
volume  method,  according  to  the  following  formula: 

md'(p-x) 

~~xd~  ~y 

m  being  the  number  of  cubic  centimeters  of  alcohol  of  density  d';  p, 
the  absolute  percentage  of  alcohol;  x,  the  desired  percentage  of  alcohol 
to  be  obtained;  d,  the  density  of  the  water  used;  y,  the  number  of 
cubic  centimeters  of  water  of  density  d  to  give  the  required  percentage 
of  alcohol.  The  formula  in  practice  was  simplified  by  taking  100  c.c. 
of  the  alcohol  and  calculating  a  table,  using  several  temperatures  as 
ordinarily  met  with  in  the  laboratory  (each  degree  from  20°  to  25°). 
The  above  formula  is  derived  in  the  following  manner: 
Let  k  =  the  absolute  weight  of  the  alcohol  taken;  then  k  =  md'p', 
where  p'  is  the  fraction  of  alcohol  in  the  absolute  alcohol  taken. 

-  (100— a;)  =  weight  of  water  to  be  added  to  make  x  per  cent  alcohol; 
whence 

-(WQ-x)=md'(l-p' 

X 


82  CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 

where  md'  (1  —  p'}  is  the  amount  of  water  in  the  alcohol  and  yd  the 
amount  of  water  to  be  added  to  make  the  amount  -  (100—  x). 

Substituting  for  k,  md'p',  the  expression  becomes 

md'p'(WQ-x)         ,,,.       „  ,      ,        rod'(100p'-oO 

-t          '-r  -      ~-   -=y 


If  p'  is  in  percentage  this  becomes 

md'(p—x)  _ 
xd        =y 

In  this  connection  a  question  arose  as  to  whether  or  not  the  volume 
method  of  making  up  the  solvent  was  experimentally  accurate.  On 
the  face  of  it  the  gravimetric  method  appears  to  be  a  safer  one,  but  it 
is  also  much  longer  and  more  tedious.  A  test  was  therefore  made  of 
both  methods.  The  calculated  amounts  of  water  and  alcohol,  that  is, 
the  apparent  weights  of  the  two  necessary  to  give  a  50  per  cent  mixture 
by  weight  (in  vacuo),  were  weighed  into  a  glass-stoppered  flask  and 
thoroughly  mixed.  A  density  determination  of  the  mixture  was  made 
at  25°  and  found  to  be  0.909826,  or  50.01  per  cent  alcohol. 

In  like  manner  a  mixture  was  made  up  by  volume,  using  the  quan- 
tities of  alcohol  and  water  calculated  from  the  above  formula  necessary 
to  make  a  50  per  cent  mixture  by  weight  ;  the  specific  gravity  in  this 
case  being  0.90980,  corresponding  to  50.02  per  cent  alcohol. 

DISSOLVED  SALTS. 

Sodium  iodide  and  potassium  iodide  were  used  in  this  investigation. 
They  were  obtained  from  Kahlbaum  and  were  extra  pure  material. 
In  fact,  it  was  not  even  necessary  for  us  to  recrystallize  them.  We 
analyzed  them,  ground  them  fine,  and  placed  them  in  a  desiccator  to 
dry  them  thoroughly  before  weighing. 

Potassium  iodide  is  not  very  soluble  in  pure  alcohol.  It  was  with 
great  difficulty  that  we  were  able  to  make  a  N/8  solution  of  it  in  the  95 
per  cent  alcohol. 

PIPETTES. 

25  c.c.,  50  c.c.,  100  c.c.,  150  c.c.,  and  200  c.c.  pipettes,  carefully 
calibrated  by  weighing  the  water  they  would  deliver,  were  employed, 
together  with  a  10  c.c.  graduated  pipette,  for  making  up  the  mixed 
solvents. 

CONDUCTIVITY  CELLS. 

The  conductivity  cells  were  of  the  same  type  as  those  used  here 
for  such  work.  The  general  method  of  determining  conductivity 
previously  described1  was  also  employed.  Since  a  number  of  changes 
in  the  temperature  regulation,  which  will  be  spoken  of  later,  were  made 
at  the  beginning  of  this  work,  and  since  these  changes  necessitated 
rewiring  of  the  system  for  the  determining  of  conductivity,  we  tested 

1Amer.  Chem.  Journ.,  46,  56  (1911). 


OF    SALTS   IN   ETHYL   ALCOHOL   AND   WATER.  83 

this  system  very  thoroughly  to  be  sure  that  all  external  resistance  in 
the  circuit  was  negligible,  or,  when  it  was  not,  we  determined  its  exact 
magnitude  in  order  to  make  the  proper  corrections. 

The  system  of  wiring  was  a  double  one,  so  that  by  means  of  a 
double-throw,  double-blade  switch,  both  a  and  b  in  the  formula  for 
calculating  conductivity 

„  va 

»>  =  KWb 

could  be  read  directly  on  the  bridge.  If  two  standard  resistance  boxes 
are  connected  one  to  each  side  of  the  bridge,  and  plugs  representing 
equal  resistances  are  removed  from  both  boxes,  then  the  reading  of  the 
wire  will  be  500  mm.,  or  exactly  its  middle-point;  provided,  of  course, 
that  there  is  no  appreciable  resistance  in  the  circuit  itself  (if  there  is  it 
must  be  evenly  balanced).  Such  was  the  case  with  our  equipment. 
Therefore,  there  were  no  corrections  of  this  kind  to  be  made  to  a  or  b 
in  the  conductivity  determinations. 

But  this  double  system  also  serves  another  purpose.  When  the  con- 
ductivity of  a  solution  is  being  measured,  if  both  a  and  b  are  read  on  the 
bridge  wire  for  the  same  resistance — supposing  that  all  other  conditions 
remain  constant — then  the  mean  of  the  two  readings  will  be  500  mm. 
If  we  do  not  find  by  actual  experiment  that  our  mean  value  is  500  mm., 
we  may  be  sure  that  at  least  one  of  the  other  conditions,  such  as,  for 
example,  temperature,  is  varying  and  needs  attention. 

It  was  very  difficult  in  some  of  the  measurements  to  obtain  distinct 
minima.  The  distances  covered  on  the  wire  on  either  side  of  the 
minimum  point  in  such  instances,  were  so  great  before  finding  corre- 
sponding sounds  on  the  two  sides  that  the  minimum  point  itself  could 
be  only  approximated.  We  endeavored  to  overcome  this  difficulty  by 
connecting  a  condenser  in  parallel  with  the  rheostat.  In  determining 
the  conductivities  of  the  alcohol  and  alcohol-water  solutions,  it  was 
practically  useless;  the  fact  is,  the  condenser  actually  made  the  read- 
ings in  some  cases  harder  to  obtain.  However,  in  determining  the 
cell  constants  of  those  cells  which  required  low  resistances,  the  bridge 
readings  were  made  much  sharper  over  a  shorter  distance,  and  the 
minimum  very  much  more  distinct  by  the  use  of  the  condenser.  We 
did  not  arrive  at  any  general  conclusion  concerning  its  use. 

TEMPERATURE  REGULATORS. 

A  number  of  different  types  of  thermo-regulators  have  been  used  in 
this  laboratory  from  time  to  time;  of  the  earlier  forms  it  is  not  necessary 
to  speak.  The  regulators  used  during  the  last  two  or  three  years  have 
had  the  general  form  of  that  shown  in  figure  31,  and  were  filled  with 
mercury. 

At  the  beginning  of  this  investigation  we  devised  a  regulator  (fig.  32) 
in  which  only  the  trap  bulbs  and  capillary  contained  mercury,  the  series 


84 


CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 


3/b' 


of  tubes  (we  used  only  two,  figure  32)  being  filled  with  toluene,  which 
has  about  6  times  the  cofficient  of  expansion  of  mercury.  Instead  of 
using  a  0.75  mm.  capillary,  into  which  a  platinum  wire  connecting  with 
the  regulating  circuit  was  introduced,  as  in  the  old  mercury  regulator, 
a  capillary  tubing  of  2  mm.  internal  bore  was  employed,  and  the  con- 
tact was  made  by  means  of  a  steel  wire  of  about  1.5  mm.  diameter  and 
rounded  at  the  end.  The  regulator  worked  with  a  fairly  high  degree 
of  accuracy,  having  a  total  variation  of  only  0.01°  at  25°  and  lower 
temperatures;  and  for  a  time  it  worked  fairly  well  at  35°,  but  it  did  not 
prove  to  be  satisfactory  for  any  length  of  time  at  the  latter  temperature. 
One  difficulty  was  that  the  mercury  at  the  surface  of  the  glass  sooner 
or  later  became  covered  with  toluene,  and  this  began  to  creep  out, 
especially  when  the  regulator  was 
not  kept  constantly  at  the  desired 
temperature,  but  was  allowed  to 
cool  down  over  night. 

This  question  in  the  meantime 
suggested  itself :  Why  is  it  not  pos- 
sible to  have  two  or  more  long 
tubes  of  thin  glass  containing  mer- 
cury instead  of  toluene,  and  in  order 
to  avoid  too  great  a  weight,  to  have 
these  tubes  of  narrow  bore?  Thus, 
we  would  have  a  greatly  increased 
surface  of  mercury,  in  comparison 
with  the  old  form  of  regulator,  and 
therefore  a  greater  expansion.  (See 
figs.  33  and  34.)  Such  a  form  was  .  . 
tried  with  great  success. 

The  supply  of  heat  to  the  ther- 
mostat was  controlled  by  a  gas- 
regulator  consisting  of  a  150-ohm  relay,  connected  electrically  with  the 
thermo-regulator  and  having  an  arrangement  attached  directly  to  the 
armature,  for  cutting  off  the  gas. 

CORRECTIONS  FOR  EXPANSION  AND  CONTRACTION. 

When  the  conductivities  of  electrolytes  in  water  as  a  solvent  are 
determined  at  the  temperatures  at  which  we  worked — that  is,  15°, 
25°,  and  35° — when  the  solutions  are  made  up  at  20°,  the  change  in 
volume  caused  by  the  expansion  or  contraction  of  the  solvent  and 
solutions  between  20°  and  these  temperatures  is  so  small  that  the 
volume  correction  can  be  neglected.  With  alcohol  and  mixtures  of 
alcohol  and  water,  however,  this  is  not  the  case.  The  expansion  here 
is  very  appreciable,  and  there  are,  in  consequence,  changes  in  the 
normality  of  the  solutions  for  which  corrections  should  be  made. 


FIGS.  31,  32,  33,  and  34. 


OF   SALTS   IN   ETHYL   ALCOHOL   AND   WATER.  85 

The  proper  corrections  have  been  applied  to  the  data  in  the  following 
tables.  The  difference  between  the  density  at  20°  and  at  the  other 
temperature  in  question  was  determined.  This  difference  represents 
the  decrease  or  increase  in  volume  per  cubic  centimeter  of  the  solution. 
Subtracting  the  decrease  below  20°  from  1.0  and  adding  it  to  1.0  above 
20°,  gives  the  coefficient  of  contraction  or  expansion  respectively. 
Since  at  15°  the  volumes  of  the  solutions  become  smaller,  there  is  a 
decrease  also  in  the  molecular  conductivities.  At  25°  and  35°  the 
expansion,  bringing  about  an  increased  volume  normality,  results  in  a 
positive  correction  to  the  conductivity. 

VISCOSITIES. 

The  viscosity  apparatus  used  in  this  investigation  was  essentially 
the  same  as  that  described  by  Davis  and  Jones1  in  their  work  on  glycerol. 
The  viscosimeters  were  of  the  general  type  therein  described,  the  capil- 
lary tubes  having  a  diameter  of  about  0.5  mm.  Some  improvements 
were  made  in  connection  with  the  constant-pressure  apparatus  for 
elevating  the  liquid  to  the  upper  mark  on  the  capillary  limb  of  the 
viscosimeter,  and  special  precautions  were  taken  to  dry  the  air  thor- 
oughly by  passing  it  through  a  long  drying-tube  filled  with  calcium 
chloride.  By  means  of  dust-traps  filled  with  cotton,  clogging  of  the 
capillary  was  effectively  prevented. 

The  desk  supporting  the  viscosimeter  stand  was  not  connected 
with  the  supports  for  the  motor  and  stirrers.  This  was  to  avoid  the 
vibration  due  to  the  motor,  and  was  secured  by  attaching  the  motor 
support  directly  to  the  walls  of  the  building.  To  reduce  the  vibra- 
tions still  further,  the  stand  holding  the  viscosimeter  rested  on  several 
layers  of  felt. 

The  stand  itself  consisted  of  a  heavy  tripodal  base,  with  a  three- 
quarter-inch  bronze  standard,  to  which  a  heavy  horizontal  arm  was 
attached  by  means  of  a  set-screw.  The  viscosimeters  were  fastened  to 
the  arm  by  means  of  a  spring  clamp,  the  tension  of  which  was  adjusted 
by  a  thumb-screw.  The  stand  was  carefully  leveled  by  means  of 
leveling  screws,  at  right  angles  to  the  line  of  sight  in  reading  the  visco- 
simeter. Leveling  in  the  other  direction  was  accomplished  by  sighting 
along  the  vertical  arm  of  the  viscosimeter  to  a  plumb-line  suspended 
before  the  glass  window  in  the  bath. 

Temperature  regulation  in  the  viscosity  work  was  essentially  the 
same  as  that  in  the  conductivity.  By  means  of  the  mercury  regulator 
(fig.  34)  already  described,  the  temperature  was  kept  constant  for  any 
desired  length  of  time  to  within  0.01°  at  15°  and  35°,  and  to  within 
0.005°  at  25°. 

JZeit.  phys.  Chem.,  81,  68  (1912). 


86 


CONDUCTIVITY   AND   VISCOSITY   OF    SOLUTIONS 


TABLE  25. — Viscosity  and  fluidity  of  potassium  iodide  in  alcohol-water  mixtures. 


Temp. 

Molecular 
concentra- 
tion. 

100  per  cent. 

95  per  cent. 

90  per  cent. 

V 

*> 

1 

•p 

•n 

<p 

15° 
25° 
35° 

fN/8  
\Solvent..  . 
fN/8 

0.01605 
0.01450 
0.01298 
0.01183 
O.OIO.'O 
0.00971 

62.32 
68.97 
77.08 
84.55 
93.46 
103.00 

0.01724 
0.01674 
0.0137J 
0.01514 
0.01098 
0.01044 

58.04 
59.70 
72.99 
76.12 
91.14 
95.81 

[Solvent..  . 
fN/8  
\Solvent..  . 



Temp. 

Molecular 
concentra- 
tion. 

80  per  cent.               70  per  cent. 

60  per  cent. 

•n 

1, 

•p                i) 

V 

15° 
25° 
35° 

/N/8  
\Solvent  .  .  . 
N/8  
Solvent... 
N/8  
Solvent.  .. 

0.02247 
0.02205 
0.01882 
0.01679 
0.01331 
0.01297 

44.49 
45.36 
53.14 
59.60 
75.13 
77.13 

0.02703 
0.02625 
0.01990 
0.01913 
0.01514 
0.01428 

37.00 
38.12 
50.25 
52.28 
66.05 
70.06 

0.03015 
0.03039 
0.02171 
0.02169 
0.01608 
0.01609 

33.16 
32.90 
46.06 
46.10 
62.22 
62.18 

Temp. 

Molecular 
concentra- 
tion. 

50  per  cent. 

40  per  cent. 

30  per  cent. 

17 

•p 

•n 

,        |         r, 

<p 

15° 
25° 
35° 

[N/8  
[Solvent  .  .  . 
fN/8  
[Solvent.  .. 
fN/8  
\Solvent.  .. 

0.03269 
0.03293 
0.02294 
0.02264 
0.01676 
0.01625 

30.59 
30.37 
43.60 
44.18 
59.67 
61.55 

0.03358 
0.03427 
0.02474 
0.02328 
0.01675 
0.01662 

29.78 
29.18 
40.43 
42.96 
59.70 
60.17 

0.03114 
0.03177 
(0.02127) 
0.02115 
0.01525 
0.01599 

32.11 

31.49 
47.12 
47.28 
65.61 
62.57 

Temp. 

Molecular 
concentra- 
tion. 

20  per  cent. 

10  per  cent. 

5  per  cent. 

r) 

V 

ri 

<p 

r) 

<p 

15° 
25° 
35° 

fN/8  
\Solvent  .  .  . 
/N/8  
[Solvent  .  .  . 

fN/8  
[Solvent  .  .  . 

0.02492 
0.02566 
0.01752 
0.01786 
0.01287 
0.01291 

40.13 
38.98 
57.09 
56.03 
77.70 
77.46 

0.01693 
0.01727 
0.01255 
0.01270 
0.00969 
0.00975 

59.06 
57.96 
79.72 
78.74 
103.2 
102.6 

0.01389 
0.01410 
0.01063 
0.01073 
0.00841 
0.00847 

72.05 
70.92 
94.07 
93.20 
118.9 
118.1 

Viscosities  at  any  temperatures  were  calculated  from  the  formula 

1)  St 

Vo  ~  s0t0 

where  t\  is  the  viscosity  coefficient  desired;  170  the  absolute  viscosity  of 
water  at  the  desired  temperatures;  s0  the  specific  gravity  of  water  *at 
that  temperature,  and  to  the  time  of  flow  of  water  in  any  given  vis- 
cometer.  s  and  t  are,  respectively,  the  density  of  the  solution  and  its 
time  of  flow  in  the  same  viscosimeter. 

The  fluidity  is  the  reciprocal  of  the  viscosity,  4>  =  - 

1 

The  percentage  temperature  coefficients  of  fluidity  are  derived  in  the 
same  manner  as  those  of  conductivity. 


OF    SALTS   IN   ETHYL   ALCOHOL   AND   WATER. 


87 


TABLE  26. — Tern 
sium 


emperature  coefficients  of  fluidity  of  polos- 
iodide  in  alcohol-water  mixtures. 


Per  cent 
alcohol. 

15  to  25° 

25  to  35° 

N/8KI 

Solvent. 

N/8KI 

Solvent. 

100.... 

95  

0.0237 

0.0226 

0.0213 

0.0217 

90.... 

0.0257 

0.0275 

0.0249 

0.0259 

80.... 

0.0194 

0.0309 

0.0338 

0.0277 

70.... 

0.0358 

0.0374 

0.0334 

0.0341 

60.... 

0.0386 

0.0401 

0.0351 

0.0349 

50.... 

0.0425 

0.0455 

0.0368 

0.0395 

40.... 

0.0358 

0.0468 

0.0476 

0.0401 

30.... 

0.0467 

0.0501 

0.0392 

0.0324 

20.... 

0.0422 

0.0437 

0.0361 

0.0382 

10.... 

0.0349 

0.0359 

0.0295 

0.0304 

5.... 

0.0306 

0.0314 

0.0265 

0.0267 

TABLE  27. — Viscosity  and  fluidity  of  sodium  iodide  in  alcohol-water  mixtures. 


Temp. 

Molecular 
concentra- 
tion: 

100  per  cent. 

95  per  cent. 

90  per  cent. 

•n 

<p 

•n 

•p 

i) 

V 

15° 
25° 
35° 

/N/8  
\Solvent  .  .  . 

/N/8 

0.01427 
0.01292 
0.01191 
0.01054 
0.00981 
0.00871 

70.21 
77.45 
83.98 
94.88 
101.9 
114.8 

0.01658 
0.01530 
0.01334 
0.01227 
0.01087 
0.00993 

60.33 
65.36 
75.01 
81.51 
92.00 
100.7 

0.01845 
0.01767 
0.01471 
0.01392 
0.01172 
0.01100 

54.22 
56.64 
67.98 
71.85 
85.34 
90.91 

[Solvent  .  .  . 

/N/8 

\Solvent  .  .  . 

Temp. 

Molecular 
concentra- 
tion. 

80  per  cent.        ;        70  per  cent. 

60  per  cent. 

•n 

¥> 

•n 

<p 

•n 

V 

15° 
25° 
35° 

/N/8  
\Solvent  .  .  . 
/N/8  
\Solvent... 
/N/8  
\Solvent.  .. 

0.02282 
0.02232 
0.01738 
0.01701 
0.01342 
0.01316 

43.78 
44.80 
57.54 
58.29 
74.52 
76.01 

0.02692 
0.02634 
0.01988 
0.01916 
0.01500 
0.01437 

37.15 
37.96 
50.31 
52.25 
66.67 
69.61 

0.03088 
0.03049 
0.02209 
0.02175 
0.01627 
0.01599 

32.38 
32.81 
45.29 
45.99 
61.48 
62.57 

Temp. 

Molecular 
concentra- 
tion. 

50  per  cent. 

40  per  cent. 

30  per  cent. 

•n 

<p 

V 

<p 

i) 

*> 

15° 
25° 
35° 

/N/8  
\Solvent  .  .  . 
/N/8  
\Solvent... 
/N/8  
(Solvent... 

0.03309 
0.03333 
0.02308 
0.02315 
0.01676 
0.01662 

30.22 
30.00 
43.34 
43.20 
59.67 
60.17 

0.03372 
0.03421 
0.02302 
0.02318 
0.01649 
0.01643 

29.65 
29.23 
43.44 
43.15 
60.66 
60.87 

0.03145 
0.03165 
0.02130 
0.02112 
0.01521 
0.01494 

31.80 
31.60 
46.95 
47.35 
65.75 
66.97 

Temp. 

Molecular 
concentra- 
tion. 

20  per  cent. 

10  per  cent. 

5  per  cent. 

i\ 

v 

•n 

V 

r> 

V 

15° 
25° 
35° 

/N/8  
\Solvent  .  .  . 
/N/8  
\Solvent  .  .  . 
/N/8  
[Solvent  .  .  . 

0.02492 
0.02551 
0.01722 
0.01766 
0.01259 
0.01296 

40.13 
39.20 
58.09 
56.67 
79.50 
77.19 

0.01739 
0.01752 
0.01288 
0.01288 
0.00987 
0.00982 

57.57 
57.09 
77.69 
77.69 
101.3 
101.8 

0.01406 
0.01410 
0.01073 
0.01072 
0.00849 
0.00847 

71.16 
90.92 
93.20 
93.28 
117.8 
118.1 

88 


CONDUCTIVITY   AND    VISCOSITY   OF   SOLUTIONS 


TABLE  28. — Conductivity  of  potassium  iodide  in  mixtures  of  ethyl  alcohol  and  water. 


Per 

F  =  N/8at20° 

Per 

F=N/128at20°. 

cent 
alco- 

15° 

25° 

35° 

alco- 

15° 

25° 

35° 

hol. 

t) 

MB 

r 

MB 

V 

MB 

V 

M» 

V 

MB 

V 

MB 

0.00 

8.000 

120.7 

8.000 

144.50 

4.98 

127.88 

90.09 

128.26 

113.37 

128.56 

138.29 

4.98 

7.993 

83.25 

8.010 

104.06 

8.035    126.30 

10.02 

127.85 

75.72 

128.28 

97.43 

128.64 

121.46 

9.55 

7.990 

70.87 

8.011 

90.44 

8.039     111.64 

20.00 

127.74 

54.231  128.39 

73.18 

128.93 

94.69 

20.00 

7.984 

49.67 

8.018 

66.46 

8.058      85.07 

21.75 

127.71 

51.061  128.41 

69.59 

128.99 

90.38 

30.06 

7.976 

38.39 

8.025 

52.78 

8.086 

69.15 

30.06 

127.61 

41.70 

128.50 

57.50 

129.25 

76.17 

40.23 

7.971 

32.34 

8.030 

44.49 

8.091 

58.44 

40.23 

127.53 

35.72 

128.58 

49.98 

129.46 

66.58 

50.61 

7.966 

30.28 

8.032 

41.07 

8.099 

53.34 

50.26 

127.49 

32.80 

128.61 

45.15 

129.56 

59.60 

60.70 

7.966 

26.49 

8.034 

35.15 

8.101 

45.00 

60.34      127.48 

30.74 

128.63 

41.39 

129.62 

53.88 

71.09 

7.960 

24.27 

8.034 

31.46 

8.101 

39.54 

70.42 

127.46 

29.85 

128.63 

39.17 

129.66 

49.52 

80.95 

7.964 

23.22 

8.035 

29.27 

8.104 

36.02 

80.51 

127.45 

29.60 

128.65 

36.92 

129.67 

46.20 

92.63 

7.962 

19.76 

8.035 

23.77 

8.104 

28.28 

90.67 

127.45 

29.05 

128.66 

35.88 

129.69 

43.48 

96.09 

7.966 

18.63 

8.035 

22.22 

8.108 

26.12 

95.77 

127.45 

28.59 

128.65 

34.61 

129.67 

40.37 

TABLE  29. — Conductivity  of  sodium  iodide  in  mixtures  of  ethyl  alcohol  and  water. 


Per 

F=N/8at20° 

Per 

F=N/32at20° 

cent 
alco- 

15° 

25° 

35° 

cent 
alco- 

15° 

25° 

35° 

V 

Mr 

V 

MB      !        " 

MB 

"            MB 

V 

MB 

V 

MB 

0  00 

8  000 

81  00 

8000 

100.4 

8.000 

121  6 

0  00 

i 

32  00 

107  40 

32  00 

130  60 

4.98 

7.993 

65.66 

8.010 

83.03 

8.035 

101.78 

4.98 

31.97 

i  69.11 

32.04 

87.75 

32.14 

108.70 

10.02 

7.991 

55.09 

8.011 

71.28 

8.040 

89.06 

10.02 

31.96 

1  58.20 

32.04 

75.61 

32.16 

95.08 

20.00 

7.984 

39.62 

8.018 

53.62 

8.058 

69.47 

20.00 

31.93 

41.90 

32.07 

57.12 

32.23 

74.50 

21.75 

7.982    i  37.48 

8.019 

51.10 

8.062 

66.54 

30.06 

31.90 

31.98 

32.10 

44.58 

32.31 

60.32 

30.06 

7.976    ]  30.82 

8.025 

42.89 

8.078 

56.69 

40.23 

31.88 

28.56 

32.12 

40.09 

32.36 

53.54 

40.23 

7.971    i  26.91 

8.030 

37.49 

8.091 

49.69 

50.26 

31.87 

25.33 

32.13 

35.04 

32.39 

47.10 

50.26 

7.968      24.66 

8.032 

33.73 

8.097 

44.21 

60.34 

31.87 

25.23 

32.13 

33.97 

32.40 

44.18 

60.34 

7.967    |  23.09 

8.033 

30.99    8.101 

39.97 

70.42 

31.86 

23.83 

32.14 

31.36 

32.42 

40.25 

70.42 

7.966      22.00 

8.034 

28.70 

8.104 

36.41 

80.51 

31.86 

24.15 

32.14 

30.71 

32.42 

38.31 

80.51 

7.966      20.78 

8.035 

26.39    8.105 

32.64 

90.67 

31.86 

23.47 

32.14 

28.69 

32.42 

34.90 

90.67 

7.966      19.42 

8.035 

23.80;  8.106 

28.63 

95.77 

31.86 

22.12 

32.14 

26.85 

32.42 

32.04 

95.77 

7.966 

18.18 

8.034 

21.88 

8.105 

25.95 

99.98 

7.966 

16.51 

8.034 

9.471 

8.103 

22.68 

Per 

F=N/128at20° 

Per 

F=N/1024at20° 

cent 
alco- 

15° 

25° 

35° 

cent 
alco- 

15° 

25° 

35° 

V 

MB 

V 

MB 

F 

MB 

.    I  » 

V 

MB 

V 

MB 

0.00 

128.00 

112.5 

128.00 

112.50 

128.00 

136.80 

0.00 

1024.0 

1024.0 

116.40 

1024.0 

141.6 

4.96 

127.88 

73.31 

128.26 

92.96 

128.56 

114.81 

5.06 

1023.0 

77.89 

1025.3 

98.87 

1028.5 

121.9 

10.00 

127.85 

61.26 

128.28 

79.83 

128.64 

100.39 

10.02 

1022.8 

65.56 

1025.4 

85.44 

1029.1 

107.6 

20.00 

127.74 

44.13 

128.39 

60.18 

128.93 

78.35 

20.00 

1021.9 

46.84 

1026.3 

63.86 

1031.5 

83.7 

29.98 

127.61 

32.83 

128.50    46.09 

129.25 

61.83 

30.06 

1020.9 

34.81 

1027.2      48.97 

1034.0 

65.8 

39.98 

127.53 

29.67 

128.58  i  41.81 

129.46 

56.02 

40.23 

1020.3 

31.72 

1027.8 

44.62 

1035.6 

60.2 

50.02 

127.49 

26.79 

128.61    34.98 

129.56 

48.04 

60.34 

1019.8 

30.38 

1028.2 

40.69 

1036.9 

53.5 

63.23 

127.48 

24.82 

128.631  36.07 

129.62 

43.42 

70.42 

1019.7 

28.51 

1028.4      37.83 

1037.3 

48.7 

70.06 

127.46 

25.99 

128.63    34.42 

129.66 

44.32 

80.51 

1019.6 

30.42 

1028.4 

39.02 

1037.4 

49.5 

80.52 

127.45 

27.07 

128.65  i  34.76 

129.67 

43.71 

90.67 

1019.6 

31.83 

1028.5 

38.39 

1037.5 

48.7 

90.61 

127.45 

27.27 

128.66    33.79 

129.69 

41.28 

95.77 

1019.6 

30.70 

1028.4 

36.46 

1037.4 

45.8 

95.00 

127.45 

26.10 

128.65 

31.88 

129.67 

38.40 

OF   SALTS   IN   ETHYL   ALCOHOL   AND   WATER.  89 

DISCUSSION  OF  THE  RESULTS. 
VISCOSITY  AND  FLUIDITY. 

Viscosity  data  have  been  obtained  in  the  various  mixtures  of  alcohol 
and  water  that  have  been  studied,  both  for  the  solvents  and  the  N/8 
solutions  of  potassium  and  sodium  iodides.  The  values  for  the  more 
dilute  solutions  approach  those  for  the  solvents  too  closely  to  be  accu- 
rately differentiated  from  them.  The  results  are  given  in  tabular  form, 
together  with  a  representative  table  of  temperature  coefficients  for  the 
range  of  temperature  over  which  the  work  was  carried  out,  i.  e.,  15°, 
25°,  and  35°. 

Table  25  gives  the  values  found  for  potassium  iodide,  and  table  27 
similar  values  for  sodium  iodide.  It  will  be  seen  that  the  effect  of  these 
salts  on  the  viscosity  of  alcohol-water  mixtures  is  comparatively  small 
for  the  N/8  solutions.  In  no  instance  does  the  decrease  in  fluidity, 
which  corresponds  to  an  increase  in  viscosity,  exceed  a  few  per  cent, 
and  in  certain  of  the  mixtures  containing  the  smaller  percentage  of 
alcohol  a  marked  increase  in  the  fluidity  is  to  be  noted.  This  will  be 
discussed  when  each  salt  is  taken  up  separately. 

In  all  mixtures  of  alcohol  and  water  from  100  per  cent  alcohol  to  that 
containing  60  per  cent  alcohol,  both  potassium  and  sodium  iodides  show 
a  marked  increase  in  the  viscosity  of  the  solvents  at  the  temperatures 
studied.  Beyond  this  point  the  effect  on  the  viscosity  is  somewhat 
different  for  each  salt. 

From  the  60  per  cent  solvent  down  to  the  0  per  cent,  i.  e.,  pure  water, 
potassium  iodide  lowers  the  viscosity  of  the  solvent  to  an  appreciable 
extent  at  15°.  At  25°  no  negative  viscosity  effect  is  to  be  noted  until  the 
30  per  cent  mixture  is  reached.  At  this  point  a  corresponding  decrease 
was  also  noted  at  35°.  Potassium  iodide,  therefore,  may  be  said  to 
increase  the  viscosity  of  all  mixtures  of  alcohol  and  water  from  100  per 
cent  alcohol  to  30  per  cent  alcohol  at  25°  and  35°,  and  to  decrease  the 
viscosity  of  all  the  other  mixtures  up  to  and  including  pure  water. 

The  shifting  of  the  point  at  which  the  fluidity  curve  for  the  salt 
crosses  that  for  the  solvent  is  to  be  accounted  for  by  the  change  in 
association  of  the  solvent  with  rise  in  temperature.  Since  a  rise  in 
temperature  causes  a  breaking  down  of  the  molecular  aggregates  of  the 
solvent  giving  ultimate  particles  of  smaller  volumes,  it  follows  that 
this  would  tend  to  shift  the  transition-point  towards  one  extreme  or  the 
other.  Since  potassium  iodide  increases  the  viscosity  of  mixtures  con- 
taining a  high  percentage  of  alcohol,  the  shifting  takes  place  towards 
the  water  end  of  the  curve. 

The  points  discussed  above  are  shown  graphically  by  figures  35  and 
36,  which  represent  the  curves  for  solvent  and  solutions  at  15°  and  25°. 
Table  26  contains  the  temperature  coefficients  of  fluidity  for  the  solvent 
and  N/8  potassium  iodide,  for  all  percentages  of  alcohol  and  water 
studied.  They  are  seen  to  decrease  in  value  with  rise  in  temperature 


90 


CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 


and  have  a  maximum  at  about  the  30  per  cent  alcohol  mixture.     The 
values  for  the  solvent  are  slightly  higher  than  those  for  the  solution. 

Sodium  iodide,  like  potassium  iodide,  increases  the  viscosity  of  mix- 
tures containing  a  higher  percentage  of  alcohol.  At  15°  the  transition- 
point  occurs  in  the  neighborhood  of  the  50  per  cent  mixture.  At  25° 
and  35°  somewhat  irregular  results  were  noted.  The  salt  increases 
the  viscosity  of  all  the  solvents  through  the  50  per  cent  mixture. 
Beyond  that  point  an  apparently  periodic  effect  occurs,  negative  vis- 
cosity appearing  only  in  the  40  per  cent  and  20  per  cent  mixtures. 
Reference  to  the  tables  will  show  that  the  difference  between  solution 

100 
95 
90 


75 

70 

£65 

a 

E  60 
55 
50 
45 
40 


75 
70 
65 
60 

55 
g 

1" 

45 


'0       10      20      30       40      50      60       70      80      90      100 
Percentage  alcohol .    x=solvent.    o=solution. 
FIG.  35. — Fluidity  of  KI  at  15°  and  25°. 


25 


0       10      20      30      40      50      60      70      80      90      100 
Percentage  alcohol.    x=solvent.    o=solution. 
FIG.  36. — Fluidity  of  Nal  at  15°. 


and  solvent  is  very  small  for  the  N/8  concentration.  In  the  10  per 
cent  and  5  per  cent  mixtures  the  salt  exerts  scarcely  any  effect  on  the 
solvent  at  25°  and  35°. 

From  this  it  would  seem  that  the  molecular  volume  of  the  dissolved 
sodium  iodide  is  smaller  than  either  the  associated  alcohol  or  water  com- 
plexes; but  in  mixtures  of  these  two  solvents  in  which  the  association 
becomes  smaller,  a  negative  viscosity  effect  is  apparent  as  soon  as  the 
dissolved  particles  are  larger  than  the  ultimate  particles  of  the  solvent. 
From  the  data  at  hand  the  change  in  association  appears  to  take  place 
more  largely  in  the  case  of  the  water  than  of  the  alcohol.  Similar 
reasoning  holds  for  potassium  iodide. 

The  explanation  of  the  phenomenon  of  negative  viscosity  as  first 
offered  by  Veazey1,  has  been  further  elaborated  by  subsequent  in- 

1Amer.  Chem.  Journ.,  37,  405  (1907) 


OF   SALTS   IN   ETHYL   ALCOHOL   AND   WATER. 


91 


vestigators1  and  it  is  unnecessary  to  discuss  it  here  in  detail.  It  is 
sufficient  to  state  that  the  facts  brought  out  in  this  investigation  are 
entirely  in  harmony  with  the  theories  established  by  previous  workers. 

CONDUCTIVITY. 

The  conductivity  data  given  above  were  plotted  in  the  form  of  curves; 
ordinates  representing  conductivities  and  abscissas  the  percentages  by 
weight  of  alcohol.  The  conductivity  of  each  concentration  of  the 
solutions  at  the  three  temperatures  were  plotted  on  one  curve  sheet 
in  order  the  better  to  compare  them.  Before  plotting  these  curves  we 
attempted  to  plot  one  which  would  give  us  the  conductivity  values 
(plotting  conductivity  against  normality)  for  the  ordinary  normalities 


j. 


10      20      30      40      50      60      70      80      90     100. 

Percentage  alcohol 
FIG.  37.— N/8-KI. 


10      20      30      40      50      60      70 

Percentage  alcohol 
FIG.  38.— N/8-NaI. 


N/8,  N/32,  etc.,  instead  of  for  those  given  above.  However,  we  found 
this  to  be  impracticable,  both  as  to  the  drawing  of  the  curves  and  also 
as  to  the  results  we  would  have  obtained;  since,  with  reference  to  the 
latter,  the  slight  change  in  conductivity  that  would  result  would  not 
be  sufficient  to  alter,  to  any  appreciable  extent,  the  character  of  the 
curves  obtained  by  using  the  above  values,  although  these  values  are 
obviously  not  strictly  comparable  with  one  another. 

It  will  be  seen  from  figures  37  and  38,  which  show  the  curves  for 
N/8  potassiu  n  and  sodium  iodides  at  15°,  25°,  and  35°,  that  there  is  a 
continual  decrease  in  the  conductivity  of  both  salts  in  passing  from 
the  pure  water  to  the  pure  alcohol. 

iCarnegie  Inst.  Wash.  Pub.  No.  80. 


92 


CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 


The  values  for  pure-water  solutions  were  taken  from  the  work  of 
West  and  Jones,1  except  those  for  15°.  For  this  temperature  they 
were  calculated  by  means  of  the  equations 


for  sodium  iodide,  and 


for  potassium  iodide.  These  equations  apply  only  to  the  N/8  solutions, 
every  other  dilution  requiring  a  different  equation.  The  accuracy  with 
which  West  and  Jones's  values  fit  the  curve  is  worth  noting. 


4.96% 


Temperature 
FIG.  39.— N/8  Nal. 


25° 

Temperature 
FIG.  40.— N/8  Nal. 


The  decrease  in  conductivity  just  mentioned  is  very  rapid  in  the 
water  end  of  the  curve  up  to  the  30  per  cent  alcohol,  and  from  there  on 
it  is  much  more  gradual.  We  may  conclude  from  this,  either  that 
the  first  addition  of  alcohol  to  water  has  a  much  more  marked  effect 
on  the  association  of  the  water  than  the  addition  of  a  little  water  to  the 
alcohol  has  on  the  association  of  the  alcohol,  or  that  comparatively 
small  quantities  of  alcohol  increase  the  viscosity  of  water  much  more 
markedly  than  is  the  case  when  small  quantities  of  water  are  added  to 
alcohol;  or,  again,  that  there  is  a  sudden  change  in  the  hydration  of  the 
dissolved  substance  caused  by  the  addition  of  the  alcohol,  whereas 
water  has  a  far  smaller  effect  on  the  alcoholation;  or,  finally,  that 
perhaps  all  or  some  of  these  factors  combine  to  produce  the  effect. 
Let  us  analyze  them. 


'Amer.  Chem.  Journ.,  34,  377  and  384  (1905). 


OF   SALTS   IN   ETHYL   ALCOHOL   AND   WATER. 


93 


The  third  conclusion  may  be  disposed  of  first  as  of  little  value,  since 
it  will  be  remembered  that  salts  of  the  alkali  metals  are  very  little 
hydrated  or  alcoholated. 

We  have  just  seen,  from  a  study  of  the  viscosity  data,  that  there  is  a 
marked  negative  viscosity,  i.  e.,  positive  fluidity,  at  the  water  end  of 
the  fluidity  curves.  On  the  other  hand,  there  is  a  steadily  increasing 
viscosity  at  the  alcohol  end.  Moreover,  the  transition-point  from 
positive  to  negative  viscosity  was  shown  to  be  nearer  the  pure  water 
than  the  pure  alcohol.  To  recapitulate,  it  was  concluded  from  these 


120 

no 
100 

90 

I  80 
|  70 
60 
60 
40 
30 


120 
IK 

100 
90 

80 

£ 

~    70 

I  60 
50 
40 


10      20      30      40      50      60      70 
Percentage  alcohol 

FIG.  41.— N/128  Nal. 


80     90     100 


20      30      40      50      60      70      80      90     100 

Percentage  alcohol 
Fio.  42.— N/32  Nal. 


results  that  the  change  in  association  brought  about  by  the  mixing  of 
the  two  solvents  appeared  to  take  place  more  largely  in  the  case  of  the 
water  than  of  the  alcohol. 

Both  our  first  and  second  conclusions,  with  regard  to  the  phenomenon 
observed  in  the  case  of  conductivity,  appear,  then,  to  be  in  perfect 
harmony  with  those  concerning  the  association  and  the  viscosity.  As 
to  the  final  conclusion,  it  obviously  follows  as  a  matter  of  course  from 
what  has  preceded,  since  the  association  and  viscosity  are  so  closely 
related. 

The  decrease  in  the  conductivity  as  a  whole  is  more  rapid  for  the 
potassium  iodide  than  for  the  sodium  iodide.  This  is  what  might  have 
been  expected  from  a  study  of  the  viscosity  data,  which,  in  turn,  are 
affected  by  the  relative  ionic  volumes  of  potassium  and  sodium.  The 
former,  having  the  larger  volume,  would  have  a  greater  negative  effect 


94 


CONDUCTIVITY   AND    VISCOSITY    OF    SOLUTIONS 


than  the  sodium  on  the  solvent,  this  being  the  same  in  both  cases. 
Although  the  ionic  velocity  of  the  potassium  is  the  greater,  this  greater 
volume  tends  to  slow  it  down  more  as  the  alcohol  is  approached. 

There  is  an  unusual  downward  deflection  in  the  curves  beyond 
the  80  per  cent  alcohol,  increasing  in  steepness  as  the  100  per  cent 
alcohol  is  approached,  which  at  first  puzzled  us  every  much.  The 
asymmetry,  that  is,  the  unsymmetrical  appearance  of  our  conductivity 
curves,  had  already  been  noted.  Then  the  question  arose,  what  would 
be  the  appearance  of  a  symmetrical  curve  if  the  viscosity  and  con- 
ductivity at  either  end  were  symmetrical  with  respect  to  the  50  per  cent 
mixture?  What  kind  of  a  curve  would  we  have?  The  answer  to  this 
question  is  given  in  figure  42,  from  which  it  can  be  seen,  on  comparison 


10     20     30     40     50     60     70     80     90     100 

Percentage  alcohol 
FIG.  43.— N/128  Nal. 


10  20  30  40  60  60  70  80  90 

Percentage  alcohol 
FIG.  44.— N/1024  Nal. 


100 


with  the  curves  for  the  actual  conductivity,  that  in  the  latter  the  point 
of  symmetry  has  been  shifted  from  the  50  per  cent  mixture  to  the  80 
per  cent,  and  that  the  end  on  the  alcohol  side  is  not  really  at  the  100 
per  cent  alcohol,  but  is  imaginary,  since  the  line  at  the  100  per  cent 
point  must  be  extended  to  make  the  curve  symmetrical.  To  be  more 
exact,  we  see  here  shifting  similar  to  that  noted  above  in  the  case  of  the 
transition-point  in  the  fluidity ;  only,  this  happens  at  every  temperature 
instead  of  with  rise  in  temperature. 

As  the  temperature  rises  the  conductivity  curves  tend  to  become 
more  and  more  nearly  a  linear  function,  that  is,  those  for  35°  have 
much  less  bend  to  them  than  those  for  15°.  A  probable  explanation 


OF   SALTS   IN   ETHYL   ALCOHOL   AND   WATER. 


95 


of  this  is  that  the  fluidities  of  a  series  of  such  mixtures  of  alcohol  and 
water  tend  also  to  become  more  nearly  a  linear  function  with  rise  in 
temperature. 

In  figure  39  we  have  plotted  the  conductivities  of  N/8  solution  of 
sodium  iodide  in  the  various  solvents  with  respect  to  temperature, 
making  ordinates  the  conductivity  and  abscissas  the  temperatures. 
From  this  we  observe  that  the 
temperature  coefficients  increase 
slightly  with  rise  in  temperature, 
with  the  exception  of  the  95  per 
cent  mixture  and  the  pure  alco- 
hol.    Further,  the  increase  be- 
comes smaller  as  we  approach  the 
pure  alcohol.     The  temperature 
coefficients  themselves  also  be- 
come smaller  as  the  percentage 
of  alcohol  becomes  greater. 

The  slight  increase  in  the  tern- 
perature  coefficients  in  the  water 
end  is  doubtless  due  to  the  same 
cause  which  produces  an  increase 
in  the  temperature  coefficients  in 
pure  water,  namely,  a  breaking 
down  of  the  hydrated  ions  with 
rise  in  temperature.  As  was  said 
before,  the  salts  of  the  alkali  metal 
are  only  very  slightly  hydrated; 
therefore,  the  small  increase  in 
temperature  coefficients. 

Further,  it  is  probable  that  since,  with  decrease  in  the  amount  of 
water,  the  increase  in  coefficients  becomes  less  until  there  is  none  in  the 
alcohol,  the  alcohol  does  not  form  alcoholates  with  these  salts. 

Thus  far  we  have  discussed  only  the  relations  which  exist  between 
the  conductivities  of  the  N/8  solutions.  Reference  to  the  curves  for 
the  other  dilutions,  figures  41,  42,  43,  44,  and  45,  will,  however,  show 
that  the  conclusions  already  arrived  at  also  hold  for  them. 

One  point  which  should  be  noted  is  that  in  the  case  of  sodium  iodide 
there  are  distinct,  though  slight  minima  in  the  curves  for  N/128  and 
N/1024  solutions,  occurring  at  about  the  70  per  cent  alcohol.  This 
phenomenon  is  not  at  all  an  unusual  one,  since  practically  everyone 
who  has  determined  the  conductivities  of  other  substances  in  alcohol- 
water  mixtures  has  found  similar  more  or  less  pronounced  minima. 
Almost  all  previous  work  was  in  dilute  solutions. 


10      20      30      40      50     60      70     80     90     100 
Percentage  alcohol 
FIG.  45. 


96  CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 

SUMMARY. 

It  has  been  our  endeavor  throughout  this  investigation  to  improve 
the  viscosity  and  conductivity  methods  wherever  possible,  in  order  to 
eliminate  the  grosser  errors  which  ordinarily  creep  in,  for  example: 
temperature  regulation;  a  more  exact  determination  of  the  external 
resistance  in  the  circuit;  a  change  from  suction  to  pressure  as  a  means 
of  raising  the  liquid  in  the  viscosimeters,  etc. 

Instead  of  making  up  our  solvents  by  the  volume  standard,  as  pre- 
vious workers  have  done,  we  used  the  weight  standard.  An  equation, 

md'(p— x) 
xd 

was  deduced  and  employed  for  determining  the  amounts  by  volume  of 
the  water  necessary  to  add  to  100  c.c.,  or  even  multiples  of  this  quan- 
tity of  alcohol,  in  order  to  make  the  required  mixtures  by  weight. 

Viscosity  and  conductivity  determinations  were  made  with  several 
dilutions  of  potassium  and  sodium  iodides,  in  a  series  of  mixtures  of 
alcohol  and  water;  and  curves  representing  fluidity  as  ordinates  and 
percentages  of  alcohol  as  abscissas  were  drawn,  as  well  as  curves  for 
conductivity  as  ordinates  plotted  both  against  mixtures  of  solvent  and 
against  temperatures  as  abscissas. 

CONCLUSIONS. 

We  arrived  at  the  following  conclusions : 

1.  The  effect  of  sodium  and  potassium  iodides  on  the  viscosity  of 
ethyl  alcohol-water  mixtures  is  comparatively  small  for  N/8  solutions. 

2.  There  is  a  marked  increase  in  the  viscosity  of  the  solvents  caused 
by  these  salts  in  passing  from  the  100  per  cent  alcohol  to  the  60  per 
cent  alcohol. 

3.  The  shifting  of  the  point  at  which  the  fluidity  curve  for  the  salt 
crosses  that  for  the  solvent  with  rise  in  temperature,  is  to  be  accounted 
for  by  the  change  in  association  of  the  solvent  with  rise  in  temperature. 

4.  This  change  in  association  is  greater  for  the  water  than  for  the 
alcohol. 

5.  The  facts  brought  out  in  connection  with  the  viscosity  work  were 
in  harmony  with  those  discovered  by  previous  workers,  and  therefore 

'can  be  explained  in  the  same  way. 

6.  There  is  a  continual  decrease  in  the  conductivity  of  N/8  sodium 
and  potassium  iodides  in  passing  from  pure  water  to  pure  alcohol.     It 
is  much  more  rapid  in  the  large  percentages  of  water  than  in  the  large 
percentages  of  alcohol. 

7.  This  may  be  explained  as  due  to  the  fact  that  the  association  of 
alcohol  is  changed  to  a  much  smaller  extent  by  adding  small  quantities 
of  water,  than  is  water  when  to  it  small  quantities  of  alcohol  are  added. 
Moreover,  since  association  and  viscosity  are  so  closely  related,  we 


OF   SALTS   IN  ETHYL  ALCOHOL   AND   WATER.  97 

also  conclude  that  the  same  reasoning  may  be  applied  to  the  latter;  in 
other  words,  there  is  a  greater  change  in  the  viscosity  of  the  water  than 
of  the  alcohol. 

8.  The  decrease  in  the  conductivity  of  potassium  iodide  with  the  in- 
crease in  the  percentage  of  alcohol,  is  more  rapid  than  the  decrease  for 
sodium  iodide,  due  no  doubt  to  the  greater  atomic  volume  of  the  former. 

9.  Hydration  has  practically  no  effect  on  the  conductivity  at  any  one 
temperature.     With  rise  in  temperature,  however,  the  breaking  down 
of  the  slightly  hydrated  ions  causes  a  small  increase  in  the  temperature 
coefficients  in  most  of  the  solutions  containing  water.     In  alcohol  the 
temperature  coefficients  are  a  linear  function,  and  therefore  there  is  no 
alcoholation. 

10.  As  the  temperature  rises  the  curves  tend  to  become  more  and 
more  nearly  a  linear  function.    We  attribute  this  to  the  fact  that  as 
the  temperature  rises  the  fluidity  curves  also  tend  to  become  more 
nearly  linear. 


CHAPTER  V. 

CONDUCTIVITY  AND  VISCOSITY  OF  SOLUTIONS  OF  RUBIDIUM  SALTS 
IN  MIXTURES  OF  ACETONE  AND  WATER. 

BY  P.  B.  DAVIS  AND  H.  HUGHES. 

The  work  done  in  the  Chemical  Laboratory  of  the  Johns  Hopkins 
University  during  the  past  dozen  years,  on  the  relations  between  the 
viscosities  of  solvents  and  solutions  of  certain  salts  in  these  solvents,  was 
referred  to  at  the  beginning  of  the  last  chapter.  That  which  is  closely 
related  to  the  contents  of  this  chapter  is  the  work  of  Jones  and  Veazey, 
Jones  and  Schmidt,  Jones  and  Guy,  and  especially  that  of  Jones  and 
Davis.1  The  last  named  extended  the  work  in  glycerol  as  a  solvent, 
studying  especially  the  conductivities  and  viscosities  of  glycerol  solu- 
tions of  ammonium  and  rubidium  salts,  as  has  already  been  pointed  out. 

AJ1  previous  work  in  the  laboratory  with  acetone  as  a  solvent  shows 
that  it  has  exceptional  properties.  Measurements  of  both  conduc- 
tivity and  fluidity  have  usually  given  results  that  are  abnormal  in 
terms  of  other  solvents:  for  this  reason  it  was  chosen  as  a  solvent  in 
this  investigation,  in  the  hope  that  the  property  possessed  by  rubidium 
salts  of  forming  solutions  having  a  lower  viscosity  than  that  of  the 
solvent  might  throw  some  additional  light  upon  the  phenomena  pre- 
sented by  solution.  Only  mixtures  of  acetone  and  water  have  been 
used,  because  in  pure  acetone  the  rubidium  salts  studied  are  not  suf- 
ficiently soluble  to  affect  the  fluidity  to  a  measurable  extent. 

EXPERIMENTAL. 

CONDUCTIVITY   APPARATUS. 

Bridge. — The  conductivity  measurements  were  made  by  means  of  a 
slide-wire  bridge  about  5  meters  long,  the  balance  being  detected  by  a 
telephone  receiver.  The  bridge  and  rheostat  were  made  and  stand- 
ardized by  Leeds  and  Northrup  Co.,  of  Philadelphia,  Pennsylvania; 
and  the  rheostat  was  compared  with  one  recently  standardized  by  the 
United  States  Bureau  of  Standards. 

A  double  system  of  wiring  was  used  between  the  rheostat,  slide  wire, 
and  cells,  so  that  by  means  of  a  double-arm  double-throw  switch  the 
arms  of  the  bridge  could  be  interchanged.  In  this  way  the  resistance 
a  of  the  first  portion  of  the  slide  wire  was  read  and  then  b  =  1000  — a 
for  comparison.  The  circuit  was  opened  and  closed  by  an  ordinary 
telegraph  key,  whose  resistance  was  made  of  negligible  value  by  con- 
necting the  frame  to  the  lever  by  a  short  spring  of  large  copper  wire. 

The  wire  used  throughout  was  number  12  gage,  and  the  cells  con- 
taining the  solutions  were  connected  with  the  rest  of  the  bridge  by  a 
large  flexible  cable  of  copper  having  a  negligible  resistance.  All  con- 
nections were  soldered,  and  the  various  portions  of  the  apparatus  were 

Carnegie  Inst.  Wash.  Pubs.  Nos.  80  and  180. 
98 


SOLUTIONS   IN   MIXTURES   OF   ACETONE   AND   WATER.          99 

tested  for  any  appreciable  resistance.  The  two  halves  of  the  double 
system  for  reading  a  and  b  were  carefully  compared,  and  b  was  found  not 
to  differ  from  1 ,000  —  a  by  any  appreciable  quantity ,  except  for  resistances 
smaller  than  any  used  in  this  investigation — that  is,  below  10  ohms. 

Cells. — The  conductivity  cells  were  of  three  forms.  For  the  most 
concentrated  solutions  two  U  cells  with  adjustable  electrodes  were 
employed,  having  constants  of  about  15,000  and  30,000.  The  most 
dilute  solutions  and  the  solvents  were  measured  in  cells  with  cylindrical 
electrodes  of  the  type  described  by  Jones  and  Schmidt,1  and  by  Jones 
and  Kreider,2  and  with  constants  ranging  from  2.9  to  4.3.  The  inter- 
mediate dilutions,  that  is,  from  N/10  to  N/400  solutions,  were  measured 
in  cells  of  the  plate  type  described  by  Jones  and  Bingham.3 

Constant-temperature  Baths. — The  constant-temperature  baths  used  in 
both  parts  of  this  investigation  were  of  the  same  general  type  employed 
for  such  work  in  this  laboratory,  and  consisted  essentially  of  round, 
galvanized-iron  tubs  of  about  20  liters  capacity,  covered  with  non- 
conducting material.  For  the  viscosity  work  the  baths  were  equipped 
with  large  glass  windows  in  the  upper  walls,  180°  apart. 

A  more  efficient  form  of  stirrer  provided  with  double  journals  and 
6-bladed  propellers  was  employed,  and  the  brackets  supporting  these 
were  attached  directly  to  the  walls  of  the  building.  The  stirrers  were 
driven  by  a  round  belt,  at  about  200  revolutions  per  minute,  by  a  -gV 
horsepower  water-cooled  hot-air  engine.  These  improvements  lessened 
materially  the  vibrations  due  to  side-thrust  from  the  propellers,  and 
increased  the  up-and-down  stirring  of  the  water  in  the  bath,  at  the  same 
time  giving  less  circular  motion. 

By  means  of  the  pressure  from  a  2.5  meter  stand-pipe,  water  could  be 
kept  flowing  through  special  cooling  coils  of  copper  placed  in  the 
bottoms  of  the  baths.  This  facilitated  temperature  regulation  at  or 
below  room  temperature.  An  auxiliary  coil  immersed  in  an  ice-bath 
was  also  introduced  into  the  cooling  system  by  means  of  brass  unions, 
whenever  the  average  temperature  of  the  tap-water  approached  too 
closely  to  the  lowest  temperature  at  which  the  work  was  attempted. 

Temperature  Regulation. — Temperature  regulation  of  a  high  degree  of 
accuracy  was  obtained  by  equipping  all  baths  with  an  approved  form 
of  electrically  operated  gas-valve,  consisting  essentially  of  a  sensitive 
150-ohm  relay,  to  the  armature  of  which  was  attached  a  device  for 
cutting  down  the  flow  of  gas  whenever  the  relay  was  set  in  action  by 
the  thermo-regulator  in  the  bath.  The  relays  were  connected  in  paral- 
lel on  a  2.5-volt  circuit  from  accumulators,  and  were  operated  by  an 
improved  form  of  mercury  thermo-regulator  of  the  general  type  de- 
scribed by  Morse,  but  having  2  to  4  reservoir  tubes  of  special  hydro- 
meter tubing,  with  walls  about  0.25  mm.  thick.  A  maximum  surface 
of  mercury  was  thus  secured,  in  keeping  with  the  stability  of  the  instru- 
ment. A  more  detailed  description  of  the  above  forms  is  to  be  found 

'Amer.  Chem.  Journ.,  42,  39  (1909).  *Ibid.,  45,  282  (1911).  »/Kd.,  34,  481  (1905) 


100  CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 

in  our  work  with  Wightman  on  conductivity  and  viscosity  in  alcohol- 
water  mixtures.1 

With  this  improved  form  of  apparatus  we  have  maintained  a  con- 
stant temperature  over  any  desired  length  of  time  to  within  0.01°,  for 
temperatures  from  15°  to  40°,  and  at  25°  with  moderate  precautions 
variations  in  temperature  not  more  than  0.005°  resulted.  An  ad- 
ditional advantage  of  the  type  of  thermo-regulators  described  above, 
over  the  more  complicated  toluol-mercury  and  other  froms,  lies  in  its 
simplicity  and  in  the  fact  that  it  may  be  readily  constructed  by  anyone 
possessing  moderate  skill  in  glass  manipulation.  The  lengths  of  the 
reservoir  tubes  need  be  limited  only  by  the  depth  of  the  baths  used. 
In  this  work  tubes  25  cm.  long  and  7  cm.  interior  diameter  were  found 
to  be  most  satisfactory.  The  thermometers  used  were  of  the  Beckman 
type  graduated  to  0.02° ;  and  these  were  compared  at  frequent  intervals 
with  a  thermometer  which  had  been  standardized  within  the  year  by 
the  United  States  Bureau  of  Standards. 

VISCOSITY   APPARATUS. 

The  viscosity  apparatus  used  throughout  this  investigation  was  the 
same  essentially  as  that  described  in  our  work  with  glycerol.2  Special 
precautions  were  taken  to  eliminate,  as  far  as  possible,  several  annoy- 
ing sources  of  error.  Vibrations  of  the  instruments  due  to  external 
causes  were  guarded  against  by  making  use  of  a  special  support  for 
the  viscosimeters,  consisting  of  a  heavy  tripodal  leveling  base  resting  on 
several  layers  of  thick  piano  felt,  and  a  large  bronze  standard  to  which 
a  horizontal  arm  was  rigidly  attached  by  a  collar  and  set-screw.  The 
viscosimeters  were  supported  in  the  bath  against  a  cork-covered  brass 
plate  at  the  extremity  of  this  arm,  by  means  of  a  spring  clamp,  the  ten- 
sion of  which  was  adjusted  to  different  instruments  by  a  thumb-screw. 

Further  precautions  were  taken  against  vibrations  by  removing  the 
engine  and  stirrer  brackets  from  direct  contact  with  the  desk  supporting 
the  baths  and  viscosimeter  stand. 

It  was  necessary  also  to  guard  against  dust  particles,  which  would 
tend  to  clog  the  capillary  of  the  viscosimeter.  To  this  end  special 
precautions  were  necessary,  both  in  making  up  solutions  and  in  using 
them  in  the  viscosimeters.  It  was  found  necessary  to  use  silk  instead 
of  linen  in  polishing  all  weighing  vessels,  and  to  wash  out  all  flasks 
with  dust-free  water  and  alcohol,  and  then  dry  them  by  a  blast  of 
air  filtered  through  cotton  wool.  The  viscosimeters  were  thoroughly 
cleansed  with  chromic  acid  before  each  procedure,  washed  as  above, 
and  dried  by  aspirating  hot,  dust-free  air  through  them.  For  this 
purpose  the  air  was  drawn  through  glass  wool,  over  calcium  chloride  in 
a  long  drying-tower,  then  through  cotton  wool,  and  finally  through  a 
short  iron  tube  heated  in  an  asbestos  chamber  by  means  of  a  flat 
burner.  A  final  filtration  through  cotton  took  place  before  the  air 

^our.  Chim.  Phys.  (1913).     Chapter  IV,  this  monograph. 
2Carnegie  Inst.  Wash.  Pub.  No.  180  (1913). 


IN   MIXTURES   OF   ACETONE   AND   WATER.  101 

was  drawn  into  the  viscosimeter.  The  instruments  were  thus  thor- 
oughly and  quickly  dried,  and  examination  with  a  hand-lens  showed 
complete  absence  of  dust  particles  in  the  capillary  or  bulb  tubes. 

When  it  was  necessary  to  take  a  series  of  readings  on  a  particular 
viscosimeter,  this  was  equipped  with  a  special  head  designed  to  exclude 
moisture  and  dust  particles  from  contact  with  the  liquid  in  the  instru- 
ment. The  liquid  was  then  raised  to  the  upper  mark  on  the  small 
bulb  by  means  of  a  constant  air-pressure  apparatus,  using  the  same 
stand-pipe  as  the  cooling  system.  Air  entering  the  viscosimeter  from 
the  pressure  vessel  was  first  carefully  dried  and  freed  from  dust  by 
the  use  of  fused  calcium  chloride  and  cotton  wool,  and  by  means  of 
stop-cocks  all  external  air  was  excluded  during  the  actual  time  of  flow  of 
the  liquid  through  the  capillary. 

By  observing  the  precautions  noted  above  we  have  succeeded  in 
obtaining  from  3  to  5  consecutive  readings  on  any  particular  viscosim- 
eter,  all  agreeing  to  within  the  limits  of  error  of  the  stop-watch  used, 
which  was  a  fine  split-second  Swiss  instrument,  reading  to  0.2  second 
and  adjusted  with  great  accuracy.  This  watch  had  the  additional 
advantage  of  running  continuously,  whether  the  hands  were  released  or 
not,  and  gave  much  better  results  than  the  intermittent  form  heretofore 
used.  Frequent  comparisons  were  made  with  standard  chronometers, 
and  no  errors  of  sufficient  magnitude  to  affect  the  accuracy  of  the  work 
were  detected. 

Specific-gravity  determinations  were  made  with  a  modified  form  of 
the  Ostwald  pycnometer,  which  is  so  well  known  that  it  does  not 
require  further  description. 

All  flasks  were  carefully  calibrated  to  hold  aliquot  parts  of  the  true 
liter  at  20°,  and  solutions  were  brought  to  within  0.1°  of  this  temperature 
before  being  diluted  to  the  mark. 

SOLVENTS. 

Water. — The  water  was  purified  by  the  method  of  Jones  and  Mackay1 
as  modified  by  Schmidt^  and  has  a  mean  specific  conductivity  of 
1.5XlO-6at25°. 

Acetone. — Kahlbaum's  so-called  pure  acetone  was  allowed  to  stand 
for  several  days  over  calcium  chloride,  and  distilled  two  or  three  times. 
No  difficulty  was  experienced  in  obtaining  a  product  of  approximately 
the  same  conductivity  as  the  water  used  in  this  work.  Solutions  were 
made  up  as  quickly  as  possible  after  distilling  the  acetone,  which  was 
always  kept  in  a  dark  place. 

Mixtures  of  Acetone  and  Water.— The  mixtures  used  as  solvents  were 
made  by  diluting  a  given  volume  of  acetone  to  a  definite  volume  with 
water  at  20°.  For  convenience,  the  number  of  cubic  centimeters  of 
acetone  diluted  to  100  c.c.  was  indicated  as  the  "percentage"  of  acetone 
in  the  solvent.  A  very  considerable  contraction  takes  place  when  ace- 
tone and  water  are  mixed,  so  that  the  figures  used  to  designate  the 

1Amer.  Chem.  Journ.,  17,  83  (1895). 


102  CONDUCTIVITY   AND   VISCOSITY   OF    SOLUTIONS 

mixture  are,  of  course,  not  true  percentages  either  by  weight  or  volume. 
The  true  percentage,  however,  is  of  no  consequence  so  long  as  the 
mixture  is  thus  denned.  Actually,  the  mixture  of  500  c.c.  acetone 
diluted  to  1  liter  at  20°  contains  about  42  per  cent  by  weight  of 
acetone,  and  the  mixture  of  250  c.c.  of  acetone  diluted  to  1  liter  con- 
tains approximately  20  per  cent  by  weight  of  acetone. 

SOLUTIONS. 

Solutions  of  one-tenth  normal  and  all  greater  concentrations  were 
made  by  dissolving  the  requisite  weight  of  salt.  The  N/50  and  N/100 
solutions  were  made  from  the  N/10  by  dilution;  and  the  N/200  and 
N/400,  respectively,  were  made  from  these.  The  N/800  and  N/1600 
concentrations  were  prepared  in  the  same  manner  from  the  N/200  and 
N/400.  The  last  two  concentrations  were  thus  made  in  three  dilutions 
from  the  N/10  concentration.  Upon  the  basis  of  a  probable  percentage 
deviation  of  0.10  per  cent  in  the  original  weight  of  salt,  and  in  the 
measurement  of  the  solvent  in  the  flask,  and  of  0.40  per  cent  in  the 
measurements  from  the  burettes,  the  probable  errors  in  the  value  of  V 
at  these  greatest  dilutions  is  about  0.70  per  cent.  The  probable  error 
in  the  other  dilutions  is,  of  course,  much  less  than  this,  being  within  0.14 
per  cent  in  the  case  of  the  N/10  and  all  greater  concentrations. 

All  solutions  were  made  up  at  20°.  No  correction  for  changed 
normality  at  higher  temperatures  had  been  applied  to  the  values 
obtained  for  molecular  conductivity.  Rise  in  temperature,  of  course, 
diminishes  the  normality  of  a  solution.  This  effect  is  accompanied  by 
an  increase  in  molecular  conductivity  which  is  complex.  While  this 
increase  in  conductivity  due  to  temperature,  is  of  the  same  order  of 
magnitude  as  that  produced  by  the  same  lowering  of  normality  caused 
by  diluting  with  more  of  the  solvent,  the  two  effects  bear  no  known 
relation  to  each  other. 

SALTS. 

The  rubidium  salts  used  in  this  work  were  Kahlbaum's  purest  pro- 
ducts. They  were  recrystallized  two  or  three  times  from  conductivity 
water,  precipitated  and  washed  with  alcohol,  and  dried  at  120°  to  135° 
according  to  the  nature  of  the  salt.  The  iodide  was  pure  white  after 
drying,  and  the  more  concentrated  solutions  were  only  slightly  colored 
after  standing  several  days. 

PROCEDURE. 

CONDUCTIVITY   MEASUREMENTS. 

The  values  of  the  molecular  conductivity  //„  are  computed  from  the 
relation  /*,  =  K  — =-,  where  v  is  the  number  of  liters  of  solvent  containing 

a  gram-molecular  weight  of  the  salt;  w,  the  resistance  in  ohms;  a/6, 
the  ratio  from  the  Wheatstone  bridge;  and  K,  the  cell  constant. 

The  cell  constants  were  determined  with  solutions  of  potassium 
chloride  of  N/50,  N/500,  and  N/200  concentrations.  The  value  taken 


IN   MIXTURES    OF   ACETONE   AND   WATER.  103 

for  the  molecular  conductivity  of  the  N/50  solution  was  that  given  by 
Ostwald,  namely,  129.7  reciprocal  Siemens  units  at  25°.  Three  resist- 
ances were  used  in  the  measurement  of  each  solution. 

It  should  be  pointed  out  that  the  fact  that  w  is  measured  in  ohms  is 
not  incongruous  with  the  use  of  reciprocal  Siemens  units  in  the  result, 
the  factor  1/1.063  being  included  in  the  cell  constants  obtained.  To 
convert  the  results  given  in  the  tables  into  reciprocal  ohms,  it  is  only 
necessary  to  multiply  by  1.063. 

The  calculation  of  data  was  facilitated  by  the  use  of  tables  of  values 

of  T  X  —  .     By  use  of  notation  by  powers  of  10,  twenty  values  of  —  were 

sufficient  for  the  preparation  of  these  tables,  which  were  readily  com- 
puted with  a  calculating  machine.  Not  only  does  the  use  of  such  tables 
save  considerable  time  in  computation,  but  in  their  preparation  the 
chances  of  mistakes  are  eliminated  by  methods  which  do  not  obtain 
when  the  values  are  independently  computed.  The  tables  used  in 
this  work  gave  all  Wheatstone-bridge  ratios  from  400/600  to  600/400 
to  millimeters  with  interpolations  to  tenths  of  millimeters. 

The  greatest  error  in  the  measurements  of  the  components  K,  v,  a,  6, 
of  the  molecular  conductivity  ju,  occurs  in  the  determination  of  the  cell 
constants  of  the  plate  cells.  This  is  subject  to  considerable  variation, 
and  it  was  necessary  to  make  frequent  determinations  of  these  values. 
Under  average  conditions  the  precision  of  n,  is  from  0.5  to  1.0  per  cent. 

Temperature  Coefficients.  —  The  temperature  coefficient  in  conduc- 
tivity units  is  the  increase  in  conductivity  for  each  degree  rise  in  temp- 
erature; that  is, 

Temp.coeff.  =^=^ 


the  temperatures  compared  always  differing  from  each  other  by  10°. 
The  temperature  coefficients  in  per  cent  is  the  above  quantity  divided 
by  the  conductivity  at  the  lower  temperature  and  multiplied  by  100, 
e.  0., 

U35°  —  M25°     100 

Percentage  coefficient  =    —^  ---  — 
Viscosity  measurements  were  calculated  from  the  formula  —  =  —  -. 

7/0  S0*0 

in  which  rj  is  the  viscosity  coefficient  for  the  liquid  in  question,  T/O  the 
absolute  viscosity  of  water,  s  the  specific  gravity  of  the  liquid  at  the 
given  temperature,  t  the  time  of  flow  of  the  same,  s0  and  t0  the  density 
and  time  of  flow  of  water  at  the  same  temperature. 

Fluidity  was  calculated  from  the  formula  <£  =  -  where  0  represents 

"n 
the  fluidity. 


104 


CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 


TABLE  30. — Molecular  conductivities  and  temperature  coefficients  of  rubidium  salts  in 
acetone-water  mixtures  at  15°,  25°,  35°,  and  45°. 

RUBIDIUM  CHLORIDE. 


Mixture. 

Molecular  conductivities. 

Temperature  coefficients. 

Conductivity  units. 

Per  cent. 

V 

15° 

25° 

35° 

45° 

15  to 

25° 

25  to 

35° 

35  to 

45° 

15  to 

25° 

25  to 

35° 

35  to 

45° 

In  75  per  cent 
acetone  

In  62.5  per  cent 
acetone  

In  50  per  cent 
acetone  

4 
10 
200 
2 
10 
200 
2 
10 
200 
2 
10 
200 
2 
10 
[  200 

f       2 

I  200 

25.5 
32.5 
50.8 
30.9 
39.0 
52.5 
39.1 
45.6 
56.7 
49.3 
54.5 
63.0 
59.0 
62.8 
73.6 
74.3 
90.5 

30.4 
43.1 
63.1 
37.7 
47.9 
66.5 
48.8 
59.5 
72.5 
61.5 

82.6 
73.9 
78.0 
94.1 
91.2 
115.1 

35.5 
50.1 
76.0 
44.8 
57.2 
82.8 
59.0 
70.9 
90.3 
74.8 
84.4 
100.1 
89.3 
93.3 
116.1 
109.4 
138.4 

40.9 

88.0 
52.8 
66.7 
97.4 

81  '.8 
108.6 
88.9 
101.5 
121.3 
106.0 
110.0 
140.0 
127.9 
164.9 

0.490 
1.063 
1.229 
0.688 
0.891 
1.20 
0.975 
1.388 
1.58 
1.225 
1.473 
1.952 
1.495 
1.520 
2.041 
1.692 
2.46 

0.502 
0.696 
1.30 
0.710 
0.922 
.632 
.020 
.144 
.78 
.325 
1.526 
.759 
.541 
.532 
2.201 
1.819 
2.33 

0.548 

i.20 
0.798 
0.955 
1.46 
.077 
.090 
.83 
.419 
.708 
.12 
.668 
.674 
2.39 
1.850 
2.65 

1.92 
3.27 
2.42 
2.23 
2.28 
2.37 
2.49 
3.04 
2.79 
2.49 
2.71 
3.10 
2.54 
2.42 
2.77 
2.28 
2.72 

1.65 
1.62 
2.06 
1.88 
1.93 
2.46 
2.09 
1.93 
2.45 
2.16 
2.21 
2.13 
2.08 
1.97 
2.34 
2.00 
2.02 

1.55 

1.58 
1.78 
1.67 
1.77 
1.82 
1.54 
2.03 
1.90 
2.03 
2.12 
1.87 
1.80 
2.06 
1.69 
1.91 

In  37.5  per  cent 
acetone  

In  25  per  cent 

In  12.5  per  cent 
acetone  

RUBIDIUM  BROMIDE. 


2 

28.2 

33.5 

38.8 

43.9 

0.529 

0.53 

0.52 

0.188 

.58 

1.34 

4 

32.5 

39.1 

45.7 

52.6 

0.655 

0.66 

0.68 

2.02 

.69 

1.49 

10 

39.1 

47.6 

55.9 

64.7 

0.845 

0.83 

0.88 

2.16 

.74 

1.57 

In  75  per  cent 

50 

48.7 

59.6 

71.0 

82.8 

1.09 

.14 

1.18 

2.25 

.91 

1.66 

aC6trOD.6 

100 

52.3 

64.2 

76.8 

89.5 

.19 

.26 

.27 

2.28 

.96 

1.65 

200 

55.1 

68.1 

81.6 

95.1 

.30 

.35 

.35 

2.36 

.98 

1.66 

400 

58.4 

71.8 

85.8 

100.9 

.34 

.40 

.51 

2.30 

.95 

1.76 

800 

61.2 

75.7 

90.8 

106.7 

.45 

.51 

.59 

2.37 

.99 

1.76 

.1600 

65.9 

81.3 

97.8 

114.9 

.54 

.65 

.71 

2.34 

.03 

1.75 

In  62.5  per  cent 

2 

36.2 

44.2 

52.7 

61.5 

.80 

.84 

.88 

2.18 

.90 

1.67 

acetone 

10 

43.5 

53.9 

63  8 

74.2 

.04 

.99 

.04 

2.39 

.84 

1.63 

200 

56.5 

72.7 

87.1 

103.9 

.62 

.44 

.68 

2.87 

.98 

1.93 

1 

40.9 

50.5 

60.6 

71.3 

.956 

.01 

.07 

2.34 

2.01 

1.76 

2 

43.3 

54.2 

65.7 

77.8 

.09 

.15 

.21 

2.51 

2.12 

1.82 

4 

45.8 

57.9 

70.6 

84.7 

.21 

.27 

.41 

2.64 

2.19 

2.00 

10 

47.8 

59.8 

72.5 

84.8 

.20 

.27 

.23 

2.51 

2.13 

1.70 

In  50  per  cent 

50 

54.6 

69.8 

86.3 

103.0 

.52 

.65 

.67 

2.78 

2.37 

1.94 

acetone  

100 

56.6 

72.4 

89.9 

107.9 

.58 

.75 

.80 

2.80 

2.42 

2.00 

200 

57.5 

73.8 

91.5 

110.0 

.63 

.77 

.85 

2.84 

2.40 

2.02 

400 

59.3 

76.1 

94.5 

113.3 

.68 

.84 

.88 

2.83 

2.42 

1.99 

800 

60.7 

78.1 

97.4 

117.2 

.74 

.93 

.98 

2.87 

2.47 

2.03 

1600 

61.6 

79.2 

98.5 

118.4 

.76 

.93 

.99 

2.86 

2.44 

2.02 

In  37.5  per  cent 

2 

51.1 

64.2 

78.4 

93.4 

.31 

.42 

.50 

2.55 

2.22 

1.92 

acetone  

10 

54.3 

67.7 

82.2 

96.5 

.35 

.55 

.43 

2.48 

2.28 

1.74 

200 

64.8 

83.8 

104.0 

125.0 

1.91 

2.02 

2.10 

2.94 

2.41 

2.02 

f       2 

61.2 

76.6 

92.6 

109.7 

In  25  per  cent 

1     10 

63.7 

79.2 

95.4 

111.6 

acetone  

1     50 

9J\J 

73.3 

93.8 

115.4 

138.4 

200 

75.7 

96.9 

119.7 

143.6 

In  12.5  per  cent 

v.  ^rVVF 
f       2 

75.6 

93.1 

111.5 

130.3 

acetone 

1   10 

81.7 

102.3 

123.5 

146.0 

IN    MIXTURES    OF    ACETONE    AND   WATER. 


105 


TABLE  30. — Molecular  conductivities  and  temperature  coefficients  of  rubidium  salts  in 
acetone-water  mixtures  at  15°,  25°,  35°,  and  46°— -Continued. 

RUBIDIUM  IODIDE.. 


Temperature  coefficients. 

TV/T     11                     rl         *•"      'f* 

Mixture. 

Conductivity  units. 

Per  cent. 

V 

15° 

25° 

35° 

45° 

15  to 

25° 

25  to 
35° 

35  to 
45° 

15  to 
25° 

25  to 
35° 

35  to 
45° 

1 

36.2 

43.0 

50.6 

57.6 

0.68 

0.75 

0.71 

1.88 

.75 

1.40 

4/3 

37.1 

44.1 

52.0 

59.8 

0.70 

0.69 

0.78 

1.89 

.56 

.50 

2 

38.8 

46.4 

53.8 

61.5 

0.75 

0.74 

0.73 

1.93 

.63 

1.36 

4 

42.7 

51.3 

59.7 

68.5 

0.87 

0.84 

0.88 

2.03 

.67 

.47 

In  75  per  cent 

10 

58.0 

68.2 

78.7 

0.93 

1.05 

.61 

.53 

acetone  

50 

67.9 

80.8 

93.7 

.30 

1.29 

.92 

.60 

100 

58.7 

71.8 

85.7 

99.5 

!38 

.39 

1.38 

2.37 

.93 

.61 

200 

60.4 

74.3 

87.6 

102.8 

.39 

.33 

1.52 

2.31 

.79 

.78 

400 

62.5 

76.3 

91.2 

107.3 

.38 

.49 

1.61 

2.21 

.95 

.77 

800 

64.2 

78.5 

94.1 

119.6 

.43 

.55 

1.55 

2.23 

.95 

.67 

1600 

65.2 

80.1 

95.3 

111.4 

.49 

.52 

1.61 

2.29 

.90 

.69 

In  62.5  per  cent 

(       2 

42.8 

53.0 

63.3 

74.2 

.01 

.13 

1.09 

2.37 

2.13 

1.73 

acetone  

10 

48.6 

60.7 

73.3 

86.7 

.23 

.26 

1.34 

2.52 

2.08 

.82 

[  200 

57.5 

72.7 

89.1 

106.5 

.52 

.65 

1.74 

2.64 

2.27 

.95 

1 

46.5 

57.4 

68.6 

81.0 

.087 

.124 

1.236 

2.34 

1.96 

1.80 

4/3 

46.7 

58.1 

69.8 

82.9 

.940 

.169 

1.314 

2.02 

2.01 

.89 

2 

47.2 

58.8 

71.0 

83.9 

.16 

.22 

1.29 

2.45 

2.08 

.82 

4 

49.0 

62.0 

76.0 

91.8 

.301 

.40 

1.575 

2.68 

2.26 

2.07 

In  50  per  cent 

10 

61.5 

65.3 

80.3 

95.2 

.379 

.502 

1.49 

2.68 

2.33 

1.86 

acetone 

50 

55.5 

71.0 

87.7 

105.7 

.55 

.67 

1.80 

2.79 

2.35 

2.05 

100 

56.6 

72.2 

108.8 

.56 

.75 

1.91 

2.76 

2.42 

2.13 

200 

58.3 

74.9 

92.8 

112.3 

.56 

.79 

1.95 

2.68 

2.39 

2.10 

400 

59.0 

75.8 

94.0 

114.1 

.68 

.82 

2.01 

2.85 

2.40 

2.14 

800 

60.2 

77.7 

95.9 

116.5 

.75 

.82 

2.06 

2.91 

2.35 

2.15 

1600 

62.5 

78.8 

96.9 

120.5 

.63 

.81 

2.36 

2.61 

2.30 

2.28 

In  37.5  per  cent 

{       2 

53.8 

67.9 

82.8 

98.7 

.41 

.49 

.59 

2.62 

2.19 

1.92 

acetone  

10 

56.2 

72.7 

89.4 

107.1 

.66 

.66 

.77 

2.95 

2.28 

1  99 

[  200 

63.8 

82.0 

101.7 

123.1 

.82 

.97 

.13 

2.86 

2.41 

2.09 

{       1 

62.7 

77.3 

92.5 

108.5 

.45 

.52 

.61 

2.32 

1.97 

1.74 

4/3 

63.0 

77.9 

94.1 

110.6 

.49 

.62 

.65 

2.36 

2.09 

1.75 

2 

63.9 

78.9 

95.7 

113.8 

.59 

.68 

.81 

2.52 

2.12 

1.89 

4 

63.0 

80.4 

98.7 

117.3 

1.65 

.84 

.85 

2.58 

2.28 

1.88 

In  25   per  cent 

10 

65.1 

83.2 

101.8 

121.6 

1.81 

.86 

.98 

2.78 

2.24 

1.94 

acetone 

50 

70.8 

90.2 

111.0 

133.4 

1.94 

2.08 

2.24 

2.74 

2.31 

2.01 

100 

72.0 

91.7 

113.1 

135.9 

1.97 

2.14 

2.28 

2.74 

2.34 

2.02 

200 

74.2 

94.5 

117.1 

139.9 

2.03 

2.26 

2.28 

2.74 

2.39 

1.95 

400 

74.0 

94.1 

116.8 

140.1 

2.01 

2.27 

2.33 

2.72 

2.42 

2.00 

800 

75.1 

95.6 

118.5 

145.2 

2.05 

2.29 

2.67 

2.73 

2.40 

2.25 

1600 

75.8 

96.4 

118.3 

141.5 

2.06 

2.19 

2.32 

2.72 

2.28 

1.96 

In  12.5  per  cent 

f       2 

76.2 

94.2 

112.5 

131.6 

1.79 

1.83 

1.92 

2.35 

1.94 

1.71 

acetone 

\    10 

80.6 

100.4 

121.4 

144.0 

1.98 

2.09 

2.26 

2.43 

2.01 

1.86 

I 

[  200 

89.9 

113.4 

137.7 

163.1 

2.35 

2.43 

2.54 

2.62 

2.14 

1.84 

106 


CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 


TABLE  30. — Molecular  conductivities  and  temperature  coefficients  of  rubidium  salts  in 
acetone-water  mixtures  at  15°,  25°,  85°,  and  45° — Continued. 

RUBIDIUM  NITRATE. 


Temperature  coefficients. 

Mixture. 

Molecular  conductivities. 

Conductivity  units. 

Per  cent. 

V 

15° 

25° 

35° 

45° 

15  to 

25° 

25  to 
35° 

35  to 

45° 

15  to 

25° 

25  to 
35° 

35  to 

45° 

4 

32.8 

39.4 

46.4 

53.7 

0.658 

0.694 

0.74 

2.01 

.77 

1.60 

10 

40.0 

48.3 

56.9 

65.7 

0.827 

0.86 

0.88 

2.07 

.78 

1.55 

In  75   per  cent 

50 

51.3 

62.5 

74.1 

86.2 

1.12 

1.16 

1.21 

2.19 

.96 

1.64 

acetone  

100 

55.4 

67.4 

80.7 

93.9 

1.20 

1.33 

1.32 

2.17 

.98 

.64 

200 

58.8 

71.6 

85.3 

99.5 

1.28 

1.37 

1.42 

2.18 

.92 

.67 

400 

61.1 

74.4 

89.0 

104.3 

1.33 

1.46 

1.53 

2.18 

.96 

.72 

800 

63.7 

79.0 

94.2 

110.6 

1.43 

1.52 

1.64 

2.21 

.93 

.74 

1600 

67.5 

82.2 

98.1 

115.3 

1.47 

1.59 

1.72 

2.18 

.94 

.75 

In  62.5  per  cent 

f       4 

38.8 

48.0 

57.8 

68.1 

0.926 

0.98 

1.03 

2.39 

2.04 

.78 

acetone  

10 

44.3 

55.2 

66.4 

78.2 

1.09 

1.13 

1.18 

2.47 

2.05 

.77 

I  200 

57.1 

71.9 

87.0 

103.8 

1.48 

1.51 

1.68 

2.59 

2.09 

.93 

2 

39.9 

50.2 

60.9 

72.4 

1.03 

1.07 

1.15 

2.59 

2.13 

.89 

4 

43.7 

55.2 

67.3 

80.7 

1.16 

1.21 

1.34 

2.65 

2.19 

.99 

10 

48.2 

61.5 

74.8 

89.3 

1.33 

.33 

1.45 

2.76 

2.16 

.94 

In  50  per  cent 

50 

54.8 

69.8 

85.8 

103.1 

1.50 

.60 

1.73 

2.74 

2.29 

2.02 

acetone  

100 

56.7 

72.4 

89.1 

107.3 

1.57 

.67 

1.82 

2.77 

2.31 

2.04 

200 

58.0 

75.6 

91.9 

110.2 

1.76 

.63 

1.83 

3.02 

2.16 

1.99 

400 

59.3 

75.5 

93.1 

112.3 

1.62 

.76 

1.92 

2.73 

2.33 

2.06 

800 

60.2 

77.1 

95.0 

115.3 

1.69 

.79 

2.03 

2.81 

2.33 

2.14 

1600 

59.2 

76.2 

94.0 

113.9 

1.70 

.78 

1.99 

2.88 

2.34 

2.12 

In  37.5  per  cent 

\      2 

45.7 

57.6 

70.5 

84.7 

1.19 

.29 

1.42 

2.61 

2.24 

2.00 

acetone 

10 

54  4 

69.4 

84.7 

101.4 

1.49 

.53 

1.67 

2.75 

2.21 

1.98 

[  200 

63.9 

81.7 

101.2 

122.0 

1.78 

.95 

2.08 

2.78 

2.39 

2.05 

2 

54.4 

68.3 

82.9 

98.2 

1.40 

.45 

1.53 

2.57 

2.13 

1.85 

4 

58.2 

73.5 

89.6 

106.9 

1.43 

.61 

1.73 

2.47 

2.19 

1.93 

10 

61.0 

75.1 

91.5 

110.3 

1.41 

.64 

1.88 

2.31 

2.19 

2.06 

In  25  per  cent 

50 

69.2 

87.8 

107.7 

128.8 

1.86 

1.99 

2.11 

2.69 

2.27 

1.96 

acetone  

100 

71.2 

90.6 

111.4 

133.5 

1.94 

2.08 

2.21 

2.78 

2.30 

1.98 

200 

73.3 

93.5 

114.9 

137.5 

2.02 

2.12 

2.28 

2.76 

2.27 

1.99 

400 

73.0 

92.9 

114.9 

137.6 

1.99 

2.20 

2.27 

2.73 

2.34 

1.98 

800 

79.4 

100.4 

123.6 

147.7 

2.10 

2.32 

2.41 

2.65 

2.31 

1.95 

1600 

73.3 

92.8 

114.2 

136.6 

1.95 

2.14 

2.24 

2.66 

2.31 

1.96 

In  12.5  per  cent 

'       2 

65.1 

80.6 

96.8 

113.7 

1.55 

1.62 

1.69 

2.38 

2.00 

1.75 

acetone  

10 

74.1 

91.7 

106.9 

123.9 

1.76 

1.51 

1.71 

2.37 

1.65 

1.60 

200 

87.0 

108.8 

131.7 

156.4 

2.18 

2.29 

2.47 

2.50 

2.10 

1.88 

IN   MIXTURES   OF   ACETONE   AND   WATER. 


107 


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CONDUCTIVITY   AND   VISCOSITY   OF    SOLUTIONS 


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IN    MIXTURES   OF   ACETONE   AND   WATER. 


109 


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CONDUCTIVITY   AND    VISCOSITY   OF   SOLUTIONS 


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IN    MIXTURES    OF   ACETONE    AND    WATER. 


Ill 


TABLE  32.— Comparison  of  the  viscosity  and  fluidity  values  for  rubidium  bromide  in 
acetone-water  mixtures. 


Temp. 

V 

75  per  cent. 

62.5  per  cent. 

50  per  cent. 

37.5  per  cent. 

•n 

<t> 

»? 

0 

•n 

* 

*7 

<*> 

15° 

f   2 
10 

0.01204 

83.07 

0.01556 
.01547 
.01516 

64.26 
64.64 
65.99 

).  01778 
.01774 
.01774 

56.25 
56.37 
56.37 

0.017SO 
.01823 
.01829 

56.19 
54.86 
54.69 

[Solv. 

.01115 

89.76 

25°  

f   2 
10 
[Solv. 

00963 

103.9 

.01200 
.01186 
.01158 

83.31 

84.29 
86.37 

.01353 
.01331 
.01329 

73.91 
75.14 
75.14 

.01331 
.01354 
.01361 

75.13 
73.87 
73.46 

00886 

112.8 

35°  

f   2 
10 

[Solv. 

00794 
66721 

125.0 
138.6 

.00951 
.00936 
.00910 

105.2 
106.8 
109.9 

.01061 
.01039 
.01026 

94.25 
96.26 
96.49 

.01028 
.01038 
.01050 

97.24 
96.34 
95.93 

45°  

f   2 
10 

[Solv. 

.00667 

149.9 

.00774 
.00757 
.00731 

129.2 
132.1 
136.8 

.00853 
.00819 
.00818 

117.2 
122.1 
122.2 

.00818 
.00823 
.00827 

122.2 
121.5 
120.9 

.00599 

166.9 

Temp. 

V 

25  per  cent. 

12.5  per  cent. 

0  per  cent. 

•n 

0 

•n 

* 

i) 

<t> 

15°  

f   2 
10 

[Solv. 

0.01630 
.01658 

61.33 
60.31 
59.05 

0.01367 
.01415 
.01430 

73.14 
70.65 
69.91 



25°  

f   2 
10 

[Solv. 

.01247 
.01227 
.01266 

80.17 
81.52 
79.01 

.01005 
.01086 
.01086 

93.87 
92.09 
92.05 

0.00872 
.00880 
.00891 

114.7 
113.6 
112.2 

35°  

f   2 
10 
[Solv. 

.00978 
.00950 
.00975 

102.3 
105.2 
102.6 

.00847 
.00852 
.00855 

118.1 
117.4 
116.9 

.00718 
.00717 
.00720 

139.4 
139.4 
138.9 

45°  

f   2 
1  10 

[Solv. 

.00790 
.00756 
.00779 

126.6 
132.1 
128.3 

.00695 
.00698 
.00698 

143.9 
143.3 
143.3 

.00608 
.00596 
.00597 

164  3 
167.7 
167.5 

DISCUSSION. 

A  parallel  investigation  of  the  viscosities  and  fluidities  of  solutions 
of  the  several  salts  studied  has  been  carried  on  in  connection  with  the 
conductivity  side  of  the  problem. 

We  have  measured  the  viscosities  of  the  tenth  and  half  normal 
solutions  of  rubidium  chloride,  bromide,  iodide,  and  nitrate,  in  all  the 
various  mixtures  of  acetone  and  water  used  as  solvents;  also  the  vis- 
cosities of  the  quarter,  three-quarters,  and  whenever  possible  the  normal 
solutions  of  the  same  salts  in  the  mixtures  designated  as  75  per  cent, 
50  per  cent,  and  25  per  cent  acetone  with  water.  It  was  not  possible 
to  obtain  data  on  solutions  of  these  salts  in  pure  acetone,  on  account 


112  CONDUCTIVITY   AND    VISCOSITY   OF   SOLUTIONS 

of  their  difficult  solubility  in  this  solvent.  However,  their  great  solu- 
bility in  water  made  it  possible  to  obtain  quite  concentrated  solutions 
in  the  majority  of  the  mixed  solvents,  even  in  that  containing  90  per 
cent  acetone.  Only  in  one  instance,  that  of  N/2  rubidium  chloride 
in  75  per  cent  acetone,  was  a  solution  obtained  which  was  non-miscible 
with  the  solvent  at  20°,  and  at  15°  a  homogeneous  solution  was  obtained. 
Table  31  contains  the  values  found. 

From  our  previous  work  on  these  salts  in  glycerol  and  water,  we 
should  naturally  expect  to  find  instances  of  negative  viscosity  in  acetone- 
water  mixtures.  However,  the  peculiarity  of  acetone  as  a  solvent  at 
once  makes  itself  evident.  Except  in  those  mixtures  containing  the 
larger  percentage  of  water,  it  will  be  noted  that  these  salts  increase  the 
viscosity  of  the  various  solvents.  Jones  and  Veazey  had  already  noted 
this  phenomenon  in  the  case  of  potassium  sulphocyanate,  but  the  nega- 
tive effect  produced  by  rubidium  salts  is  so  great  in  other  solvents  that 
the  two  classes  of  salts  can  hardly  be  regarded  as  comparable. 

A  glance  at  the  tables  will  show  that  rubidium  iodide  and  nitrate,  the 
two  salts  found  to  give  the  greatest  viscosity  lowering  in  glycerol- water 
and  their  mixtures,  produce  a  marked  increase  at  all  dilutions  in  the 
viscosity  of  the  solvents  up  to  the  50  per  cent  acetone  mixture.  Beyond 
this  point  the  fluidity  curve  (fig.  46)  for  the  salts  crosses  that  of  the 
solvent,  and  a  negative  viscosity  effect  becomes  apparent  in  the  mix- 
tures containing  the  lower  percentages  of  acetone.  The  50  per  cent 
mixture  is  apparently  very  close  to  the  transition-point,  since  certain 
dilutions  apparently  increase  the  viscosity  of  the  solvent,  while  others 
lower  it.  It  would  seem  that  in  mixtures  from  100  per  cent  to  50  per 
cent  acetone  the  molecular  volume  of  the  dissolved  salt  is  smaller  than 
the  molecular  aggregates  of  the  solvents;  and  in  the  other  mixtures, 
larger.  The  salts  lower  the  viscosity  of  pure  water,  because  according 
to  Jones  and  Veazey's  theory  their  molecular  volumes  are  greater  than 
the  complexes  of  the  solvent.  On  the  addition  of  acetone  having  appar- 
ently much  larger  molecular  complexes,  this  negative  viscosity  effect 
becomes  less  and  less  with  increasing  percentage  of  acetone,  until  we 
reach  a  mixture  in  which  the  two  factors  balance  one  another.  This 
point  is  in  the  neighborhood  of  the  50  per  cent  mixture.  By  still 
further  increasing  the  percentage  of  acetone,  the  aggregates  of  the 
solvent  exceed  the  molecules  and  ions  of  the  solute  in  size  and  a 
positive  viscosity  effect  results. 

Associated  with  each  table  of  viscosities  and  fluidities  is  a  corre- 
sponding table  of  temperature  coefficients.  Their  relations  to  those  of 
conductivity  are  taken  up  in  the  discussion  of  that  phase  of  the  work, 
which  immediately  follows. 

It  was  found  by  Jones  and  Veazey1  that  the  curves  expressing  the 
fluidity  of  varying  mixtures  of  acetone  and  water  are  almost  exactly 

.  Chem.  Journ.,  37,  405  (1907). 


IN   MIXTURES   OF   ACETONE   AND   WATER. 


113 


parallel  to  those  for  the  conductivity  of  potassium  sulphocyanate  in 
solution  in  the  same  mixtures.  The  fluidity  curve  for  acetone  has  a 
minimum  between  37.5  per  cent  and  50  per  cent;  and  the  fluidity  for 
the  rubidium  halides  and  nitrate  has  its  minimum  in  approximately  the 
same  position  (fig.  46).  The  conductivity  curves,  however,  of  the  rubi- 
dium salts  have  their  minima  corresponding  to  a  much  greater  per- 
centage of  acetone  (fig.  47).  As  has  been  shown  by  previous  workers, 
this  minimum  in  fluidity  occurs  at  the  position  where  the  breaking 
down  of  association  of  one  solvent  by  the  other  is  greatest.  The  con- 
ductivity depends  upon  the  velocities  of  the  ions  and  the  degree  of 
dissociation.  The  dissociation  is  least  when  the  association  of  the 
solvent  is  least,  and  the  speed  of  the  ions  is  least  when  the  fluidity  is 
greatest.  Therefore,  if  these  were  the  only  determining  factors,  the 
conductivity  minima  always  correspond  to  the  fluidity  minima. 


240 
22L) 
200 
180 
160 
140 
120 
100 

80 

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100      87.5       75       62.5       50       37.5 
Percentage  acetone 

FIG.  46. — Fluidity  of  rubidium  bromide, 
solution  and  solvent. 


, 

/ 

/ 

7 

,, 

"/ 

/ 

\ 

.          / 

/ 

^  — 

^ 

/ 

/ 

75          62.5          50          37.5          25           12.5           0 

Percentage  acetone 

FIG.  47. — Conductivity  and  viscosity  of  rubid- 
ium iodide  in  acetone-water  at  25°. 
Curve  I,  ordinates,  molecular  conductivity. 
Curve  II,  ordinates,  fluidity. 


Potassium  sulphocyanate  has  a  considerable  solubility  in  pure  ace- 
tone (about  20  grams  in  100  grams  of  acetone  at  20°),  whereas  the 
rubidium  salts  studied  are  only  slightly  soluble  in  pure  acetone.  There- 
fore, in  the  same  concentrations,  the  rubidium  salts  are  nearer  satura- 
tion than  potassium  sulphocyanate.  The  percentage  dissociation  is, 
therefore,  lower  in  the  case  of  rubidium  salts  than  in  that  of  potassium 
sulphocyanate.  A  possible  explanation  of  the  shifting  of  the  minimum 
in  the  conductivity  of  rubidium  salts  towards  the  greater  proportions  of 
acetone,  is  that  the  great  insolubility  in  acetone  might  cause  the  disso- 


114 


CONDUCTIVITY   AND   VISCOSITY   OF   SOLUTIONS 


elation  to  be  driven  back.  This  shifting  of  the  minimum  by  the  slight 
solubility,  however,  seems  to  be  clearly  manifested  only  between  salts 
with  great  difference  in  solubility  in  acetone,  as  is  the  case  with  potas- 
sium sulphocyanate  and  the  rubidium  halides.  Of  the  four  salts,  rubi- 
dium chloride,  iodide,  bromide,  and  nitrate,  the  solubility  in  acetone  of 
only  the  iodide  is  accurately  known,  so  that  these  salts  could  not  be  com- 
pared with  each  other  for  the  relation  between  solubility  and  minimum 
conductivity. 

The  relative  solubilities  in  water  probably  do  not  correspond  to  the 
relative  solubilities  in  acetone.  The  iodide  and  the  nitrate  are  respec- 
tively the  most  soluble  and  the  least  soluble  in  water;  but  the  conduc- 
tivity minima  of  these  two  salts  are  the  farthest  from  the  ordinate, 
corresponding  to  100  per  cent  acetone  (figs.  48  to  51). 


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/ 

/ 

N 
10 

/ 

/ 

1 

Percentage  acetone 

FIG.  48. — Conductivity   of   rubidium 
iodide  in  acetone-water  at  25°. 


75    62.5     50      37.5     25     12.50     0 
Percentage  acetone. 

FIG.  49. — Conductivity   of   rubidium 
bromide  in  acetone-water  at  25°. 


It  is  not  to  be  expected  that  the  differences  in  solubility  in  water  of 
the  different  rubidium  salts  would  show  this  relation,  because  in  the 
dilutions  which  are  sufficiently  great  to  give  any  minimum  at  all,  these 
salts  are  very  far  from  their  saturation  in  water.  In  the  dilutions 
greater  than  N/800,  that  is,  where  the  dissociation  approaches  com- 
pleteness, the  minima  in  the  conductivity  curves  are  seen  to  be  nearer 
those  of  fluidity  (figs.  48  and  51). 

A  comparison  of  the  percentage  coefficients  of  fluidity  given  in 
tables  34  to  38  with  those  of  conductivity,  shows  that  the  two  are 
nearly  equal,  which  is  to  be  expected  in  the  case  of  a  non-solvated  salt. 


IN   MIXTURES   OF   ACETONE   AND   WATER. 


115 


For  the  same  salt  the  coefficients  of  conductivity,  while  nearly  equal, 
are  somewhat  smaller  than  those  of  fluidity.  As  Davis  and  Jones  have 
pointed  out,  this  is  due  to  "the  decrease  in  association  of  the  solvent 
with  rise  in  temperature,  causing  a  decrease  in  the  ionization  of  the 
solute,  and  therefore  a  smaller  conductivity." 

In  most  cases  the  molecular  conductivities  of  N/800  and  N/1600  are 
practically  the  same,  showing  that  the  dissociation  has  apparently 
become  nearly  constant.  If  the  conductivity  depends  only  upon  the 
velocity  of  the  ions  and  the  number  of  ions  present,  then,  in  the  case 
of  a  non-hydrated  electrolyte,  since  velocity  is  proportional  to  fluidity, 
fj.^  in  solvent  <p  of  solvent 


in  water 


(p  of  water 


/ 

/ 

/ 

/ 

/ 

/ 

/ 

/ 

/I 

/ 

-r 

200 

/ 

/ 

/ 

/ 

/ 

4 

10 

/ 

iL 

75    62.5     50     37.5      25      12.5      0 

Percentage  acetone. 

FIG.   50. — Conductivity   of  rubidium 
chloride  in  acetone-water  at  25°. 


75     62.5     50     37.5     25      12.5      0 
Percentage  acetone. 

FIG.  51. — Conductivity  of   rubidium 
nitrate  in  acetone-water  at  25°. 


This  would  indicate  a  value  of  /*«>  =  102  for  rubidium  bromide  at  25° 
in  "50  per  cent"  acetone.  The  molecular  conductivity  becomes  con- 
stant at  about  80.  This  indicates  either  that  the  equilibrium  between 
ions  and  molecules  becomes  constant  at  a  =  80/102  =  78  per  cent 
(where  a  is  the  percentage  dissociation) ;  or  that  the  dissociation  is 
complete  at  the  N/800  dilution,  and  the  molecular  conductivity  is 
decreased  by  a  decreasing  velocity  of  the  ions.  The  first  alternative 
seems  improbable.  The  second  may  seem  unlikely  in  view  of  the  fact 
that  at  high  concentrations  rubidium  salts  are  not  appreciably  solvated. 
This,  however,  is  not  evidence  that  there  is  no  solvation  at  great  dilu- 
tion. And  this  would  seem  to  be  the  most  probable  explanation  of  the 
low  constant  value  for  /*. 


116  CONDUCTIVITY   AND    VISCOSITY   OF    SOLUTIONS. 

SUMMARY. 

The  viscosities  and  conductivities  of  a  number  of  rubidium  salts  have 
been  measured  in  various  mixtures  of  acetone  and  water. 

Rubidium  salts  increase  the  viscosity  of  all  mixtures  containing  a 
larger  percentage  of  acetone. 

The  curve  representing  the  fluidities  of  a  solution  of  any  of  these 
salts  in  the  various  solvents  crosses  the  curve  for  the  solvents  in  the 
neighborhood  of  the  50  per  cent  acetone-water  mixture. 

Negative  viscosity  coefficients,  wherever  found,  were  much  smaller 
than  corresponding  values  in  water  or  glycerol. 

The  temperature  coefficients  of  fluidity  of  acetone-water  mixtures 
are  very  small  and  decrease  with  rise  in  temperature. 

The  largest  temperature  coefficients  occur  in  the  mixture  containing 
50  per  cent  acetone,  i.  e.,  the  curve  representing  temperature  coeffi- 
cients passes  through  a  maximum  at  that  point. 

Minima  in  the  conductivity  curves  for  rubidium  salts  correspond  to 
a  higher  percentage  of  acetone  in  the  solvent  mixtures  than  do  those 
in  the  fluidity  curves,  whereas  the  two  curves  are  parallel  for  certain 
other  salts.  A  possible  explanation  based  on  the  difference  in  solu- 
bility is  offered. 

A  comparison  of  the  temperature  coefficients  of  conductivity  and 
fluidity  shows  that  these  are  what  is  to  be  expected  in  the  case  of  a 
non-solvated  salt  in  a  mixture  of  associated  solvents. 

A  possible  indication  of  solvation  of  rubidium  salts  in  dilute  solutions 
is  pointed  out. 


CHAPTER  VI. 

THE  CONDUCTIVITY  AND  VISCOSITY  OF  CERTAIN  RUBIDIUM  AND 

AMMONIUM  SALTS  IN  TERNARY  MIXTURES  OF  GLYCEROL, 

ACETONE,  AND  WATER  AT  15°,  25°,  AND  35°. 

BY  P.  B.  DAVIS  AND  W.  S.  PUTNAM. 
INTRODUCTION. 

The  fairly  extensive  investigations  of  Jones  and  his  collaborators  on 
conductivity  and  viscosity  in  the  field  of  mixed  solvents,  have  been 
brought  together  and  correlated  in  two  elaborate  monographs  pub- 
lished by  the  Carnegie  Institution  of  Washington.1 

By  far  the  greater  part  of  this  previous  work  has  been  devoted  to 
binary  mixtures  of  the  various  solvents  studied,  as  well  as  to  the  pure 
solvents  themselves.  Up  to  the  present,  the  work  has  covered  very  thor- 
oughly the  determination  of  the  conductivity  and  viscosity  coefficients 
of  a  large  number  of  compounds,  both  inorganic  and  organic,  in  water, 
and  in  acetone,  glycerol,  and  the  alcohols,  as  well  as  in  binary  mixtures 
of  the  latter  solvents  with  one  another  and  with  water.  Thus  far,  how- 
ever, few  if  any  attempts  have  been  made  to  carry  out  a  systematic 
study  of  the  behavior  of  such  compounds  in  ternary  mixtures  contain- 
ing the  above-named  solvents.  Such,  then,  has  been  the  object  of  the 
present  investigation,  which  may  be  taken  as  the  initial  step  in  a  series 
of  similar  researches. 

Before  taking  up  the  discussion  of  this  phase  of  the  subject,  a  short 
review  of  the  various  relations  and  deductions  brought  out  by  previous 
investigators  in  the  field  of  mixed  solvents,  should  serve  as  a  fitting 
introduction  to  the  present  work,  by  calling  to  mind  the  various  lines 
of  evidence  bearing  on  this  subject. 

However,  since  we  have  been  concerned  more  particularly  with 
glycerol,  acetone,  and  water  in  this  and  in  previous  contributions  to  the 
literature  on  the  subject,  the  review  following  will  be  confined  to  the 
investigations  covering  these  three  important  solvents.  Moreover,  the 
work  in  mixed  solvents  containing  the  alcohols  has  recently  been  care- 
fully reviewed  in  a  previous  article. 

The  first  important  work  in  mixed  solvents  contai  ning  acetone  was 
that  of  Jones  and  Veazey.  Prior  to  their  investigations,  Bingham  and 
others  had  made  some  preliminary  determinations  of  conductivities 
and  fluidities  in  this  solvent  along  with  their  work  in  the  alcohols. 
Thus,  Bingham  noted  the  characteristic  minima  occurring  in  the  con- 
ductivity curves  for  certain  salts  in  acetone-water  mixtures,  and  pointed 
out  that  a  connection  undoubtedly  existed  between  this  and  a  similar 
phenomenon  in  the  fluidity  curves  for  such  mixtures. 

Carnegie  Inst.  Wash.  Pubs.  Nos.  80  and  180. 

117 


118  CONDUCTIVITY   AND   VISCOSITY 

Subsequently,  McMaster  found  these  minima  to  be  more  pronounced 
at  lower  temperatures,  and  corroborated  the  observations  of  Bingham 
regarding  the  relations  between  conductivity  and  fluidity  minima  in 
these  curves.  He  also  noted  and  offered  a  tentative  explanation  of 
certain  maxima  in  the  curves  for  acetone-alcohol  mixtures. 

Although  Veazey's  work  has  been  fully  reviewed  in  a  previous 
article,  it  bears  directly  on  the  present  investigation,  since  it  contains 
some  of  the  facts  earlier  established.  In  addition  to  confirming  the 
above-mentioned  deductions  of  Bingham,  McMaster,  and  others, 
Veazey  noted  and  explained  the  marked  increase  in  viscosity  on  mixing 
acetone,  as  well  as  the  alcohols,  with  water.  This  he  showed  to  be  due 
to  a  mutual  diminution  in  the  association  of  the  respective  solvents,  the 
resulting  mixture  having  a  greater  number  of  ultimate  particles  and 
hence  a  larger  viscosity  coefficient  than  either  solvent  separately.  More- 
over, Veazey  was  the  first  to  offer  an  entirely  satisfactory  explanation 
of  the  phenomenon  of  negative  viscosity  noted  in  certain  aqueous 
solutions  by  a  number  of  previous  investigators,  and  extended  this  field 
to  include  mixed  solvents.  His  interpretation  of  this  phenomenon  is 
now  too  well  known  to  require  more  than  the  mere  statement  that  it  is 
based  on  the  relations  of  the  molecular  volume  of  the  solute  to  that  of 
the  solvent,  negative  viscosity  occurring  only  when  the  former  is  much 
greater  than  the  latter.  This  relation  is,  furthermore,  borne  out  by  the 
position  of  the  cations  of  the  solutes  causing  negative  viscosity  at  the 
maxima  of  the  atomic  volume  curve  of  Lothar  Meyer. 

Jones  and  Schmidt  studied  glycerol  as  a  solvent,  and  carried  out 
determinations  in  both  the  pure  and  mixed  solvents.  They  found  it 
well  adapted  to  both  conductivity  and  viscosity  work,  since,  in  addi- 
tion to  possessing  a  high  viscosity  coefficient,  it  proved  to  be  a  good  dis- 
sociant,  and  showed  the  largest  temperature  coefficients  of  conduc- 
tivity and  viscosity  of  any  solvent  hitherto  employed. 

Guy  and  Jones  extended  greatly  the  field  opened  up  by  Schmidt,  and 
from  a  large  number  of  measurements  pointed  out  that  molecular  con- 
ductivities in  glycerol  are  extremely  small,  but  show  a  regular  increase 
with  dilution  and  rise  in  temperature.  It  was  also  shown  that  salts 
having  the  greatest  hydrating  power  in  water  possess  the  largest  tem- 
perature coefficients  of  conductivity  in  glycerol.  In  mixed  solvents 
Guy  and  Jones  found  that  conductivities  do  not  follow  the  law  of  aver- 
ages, but  are  always  smaller,  and  that  the  ternary  electrolytes  produce 
a  greater  increase  in  the  viscosity  of  the  solvent  than  the  binary  elec- 
trolytes. Isolated  instances  of  negative  viscosity  were  observed  both 
in  glycerol  and  in  certain  mixtures  of  glycerol  and  water,  which  led 
Davis  and  Jones  to  make  a  closer  study  of  this  phenomenon. 

Davis  and  Jones,  working  from  the  standpoint  of  negative  viscosity, 
made  a  careful  study  of  the  conduct  of  rubidium  and  ammonium  salts 
both  in  glycerol  and  in  glycerol-water  mixtures.  They  found  that 


OF   CERTAIN   SALTS   IN   TERNARY   MIXTURES.  119 

rubidium  salts  produced  a  phenomenal  lowering  of  the  viscosity  of 
glycerol,  while  ammonium  salts  proved  to  be  more  closely  allied  to 
rubidium  than  to  potassium  in  their  effects  on  a  solvent  like  glycerol. 
They  also  noted  minima  in  certain  of  the  conductivity  curves  for  the 
more  concentrated  solutions  studied;  the  conductivity  varying  directly 
with  the  fluidity.  In  addition  to  this  the  percentage  increase  in  fluidity 
was  found  to  diminish  rapidly  with  rise  in  temperature  and  with  dilu- 
tion, and  the  curves  representing  fluidity  and  conductivity  in  glycerol- 
water  mixtures  showed  marked  similarity.  No  evidence  of  positive 
viscosity  of  solutions  of  rubidium  salts  in  glycerol  was  found,  and  in 
the  case  of  mixed  solvents  only  at  comparatively  high  temperatures. 

The  study  of  the  behavior  of  rubidium  salts  in  mixed  solvents  was 
extended  by  Davis,  Hughes,  and  Jones  to  acetone-water  mixtures.  A 
marked  increase  in  viscosity  was  found  for  all  the  rubidium  salts  in  the 
solvents  containing  the  larger  percentage  of  acetone,  a  phenomenon 
which  this  electrolyte  had  exhibited  in  none  of  the  solvents  previously 
studied.  The  curve  representing  the  fluidity  of  solutions  of  these  salts 
in  the  different  mixtures  was  observed  to  cross  that  of  the  solvents  in 
the  vicinity  of  the  mixture  containing  50  per  cent  acetone.  Negative 
viscosity  coefficients,  wherever  noted,  were  much  smaller  than  corre- 
sponding values  in  glycerol-water  mixtures.  A  comparison  of  the  tem- 
perature coefficients  of  fluidity  and  conductivity  showed  them  to  be 
very  similar,  and  of  the  order  of  magnitude  to  be  expected  for  a  non- 
solvated  salt  in  a  mixture  of  associated  solvents.  In  addition,  minima 
were  noted  in  the  conductivity  curves  for  these  salts,  and  were  found  to 
correspond  to  a  higher  percentage  of  acetone  in  the  solvent  than  in  the 
case  of  similar  minima  in  the  fluidity  curves. 

The  important  observations  on  solutions  in  binary  mixtures  made  by 
the  above-mentioned  investigators  in  this  laboratory  make  it  evident 
that  some  lines  of  evidence  are  desirable,  on  the  behavior  of  certain 
salts  in  ternary  mixtures  containing  the  three  important  solvents  dis- 
cussed, viz,  glycerol,  acetone,  and  water.  The  present  investigation, 
therefore,  has  been  devoted  to  a  study  of  the  behavior  of  rubidium  and 
ammonium  salts  which  exhibit  negative  viscosity  to  a  high  degree  in 
many  pure  solvents  and  their  binary  mixtures,  in  a  new  series  of  sol- 
vents which  contain  varying  proportions  of  glycerol,  acetone,  and  water. 

EXPERIMENTAL. 
APPARATUS. 

Thermostats. — As  in  previous  years,  it  has  been  our  constant  aim  to 
bring  to  as  near  perfection  as  possible  this  fundamental  part  of  the 
apparatus.  With  this  in  view,  a  new  type  of  thermostat  (fig.  52  A) 
has  been  devised  suitable  both  for  conductivity  and  viscosity  determi- 
nations, or  for  reaction  velocity  work;  and  three  such  baths  have  been 
recently  installed  and  put  into  full  working  order  (Plate  I).  These 
thermostats  are  of  about  60  liters  capacity  and  are  substantially  con- 


120  CONDUCTIVITY   AND   VISCOSITY 

structed  of  copper.  Gas  is  employed  as  a  means  of  maintaining  the 
desired  temperature,  the  heat  being  applied  to  a  heavy  iron  pipe  (fig. 
52  A-H)  outside  the  circumference  of  the  bath,  and  through  this  pipe 
water  is  kept  circulating  by  the  propellers.  Thus  only  a  small  portion 
of  the  water  in  the  thermostat  comes  into  immediate  contact  with  the 
heated  surface,  being  subsequently  mixed  with  the  main  body  of  water, 
thereby  securing  much  more  even  distribution  of  heat. 

The  improved  type  of  mercury  regulator  described  by  Davis  and 
Hughes  was  used  to  operate  the  relays  (fig.  52  c)  controlling  the  gas- 
supply  to  the  micro-burners  of  the  thermostats.  A  new  type  of  toluene 
regulator  suitable  for  a  wide  range  of  temperature  has  been  constructed 
(fig.  52  B).  With  the  above  improved  apparatus,  temperature  regu- 
lation to  within  0.01°  was  easily  maintained  over  any  length  of  time 
throughout  the  work,  and  with  a  reasonable  amount  of  attention  regu- 
lation to  0.005°  was  attained.  (See  fig.  52  for  details  of  the  system.) 

The  hot-air  engines  formerly  used  as  a  source  of  power  for  the  stirrers 
have  been  discarded  in  favor  of  the  electric  motor,  which  gives  greater 
freedom  from  vibrations  and  permits  the  maintainence  of  constant 
temperature  in  the  baths  both  day  and  night.  Both  belt  and  friction 
drive  were  used  to  transmit  power  from  motor  to  stirrers,  a  1-12  horse- 
power direct-current  motor  serving  to  operate  all  five  thermostats. 

Conductivity  Apparatus. — The  conductivity  apparatus  used  in  this 
investigation  was  identical  with  that  employed  for  our  recent  work  in 
acetone-water  mixtures ;  the  methods  of  obtaining  duplicate  readings, 
system  of  wiring,  and  similar  details  remaining  exactly  the  same  as  in 
the  earlier  work.  The  conductivity  cells  also  were  of  the  type  now 
generally  employed  here  for  such  work,  and  have  been  fully  described 
elsewhere.  All  the  instruments  used  were  carefully  calibrated  at  regu- 
lar intervals  or  compared  with  standards. 

Viscosity  Apparatus. — The  usual  type  of  viscosity  apparatus  as 
developed  and  improved  here  was  used  throughout  this  work.  Special 
care  was  exercised  both  in  calibrating  all  the  instruments  and  in  guard- 
ing against  external  jars  and  vibrations. 

Volumetric  Apparatus. — All  flasks  were  carefully  calibrated  by 
repeated  weighings  to  hold  aliquot  parts  of  the  true  liter  at  20°;  and 
solvents  and  solutions  were  brought  to  within  0.1°  of  this  temperature 
before  being  diluted  to  the  calibration  mark.  All  pipettes  used  in 
making  up  the  solvents  were  standardized  to  drain  a  definite  amount 
of  each  component  at  20°,  the  mean  of  a  number  of  weighings  being 
taken  as  the  drainage  capacity  of  the  respective  instruments. 

SOLVENTS. 

The  glycerol  used  in  these  ternary  mixtures  was  Kahlbaum's  1.26, 
with  a  mean  specific  gravity  of  1.257  at  25°  and  a  specific  conductivity 
of  0.6X10~7  at  the  same  temperature. 


Enlarged  View  of  New  Form  of  Thermostat.      For  diagram  and  detailed  description  see  Fig.  52. 


General  View  of  Constant  Temperature  Apparatus   for  Viscosity  and  Conductivity   Investigations 
showing  arrangement,  metnods  of  stirring,  heating,  regulating,  etc. 


OF   CERTAIN    SALTS   IN   TERNARY   MIXTURES. 


121 


FIG    52  -Constant  Temperature  Apparatus  for  Conductivity  and  Viscosity  Investigations. 


and  rougl1 


filled  with  waste;  (BE)  oil  trap;  (FF)  six-bladed 

E-  ^inss  ss  rs»? 

very'low  or  high  temperatures);    ™coolmg  coils;  (oo) 


in  cap-   (KK)   regulating  valve  to 

jt  s  w  ta  ^-^  ^ 

„„,  „  to  E.  (pp)  aper. 


122  CONDUCTICITY   AND   VISCOSITY 

The  acetone  was  dried  over  calcium  chloride  for  at  least  a  week  before 
using  and  was  then  redistilled  several  times  immediately  before  making 
up  the  mixed  solvents.  It  had  a  mean  specific  gravity  of  0.787  at  25°, 
and  a  specific  conductivity  of  about  4Xl(T7  at  that  temperature. 

The  conductivity  water  was  obtained  by  the  method  of  Jones, 
McKay,  and  Schmidt,  and  had  a  mean  specific  conductivity  of 
1.5XKT6  at  25°. 

The  mixed  solvents  were  prepared  by  mixing  one  or  two  parts  of  each 
of  the  above  components  with  varying  proportions  of  the  other  two. 
Seven  such  combinations  proved  to  be  possible,  and  were  prepared  for 
solvents  in  one-liter  quantities  immediately  before  using. 

The  specific  data  relating  to  each  solvent  are  to  be  found  in  table  33. 


The  rubidium  and  ammonium  salts  were  all  carefully  recrystallized 
from  conductivity  water,  precipitated  and  washed  with  absolute 
alcohol,  then  dried  first  in  the  steam  oven,  and  finally  pulverized  and 
heated  in  an  air-bath  at  the  most  favorable  temperature  for  the  salt 
in  question.  By  this  procedure,  products  of  an  exceptional  purity  were 
obtained,  and  even  in  the  case  of  ammonium  iodide  the  concentrated 
solutions  became  only  slightly  tinted  after  standing. 

SOLUTIONS. 

All  solutions  were  made  up  as  described  by  Davis  and  Hughes;  the 
concentrated  solutions  by  direct  weighing,  the  others  by  successive 
dilutions.  All  operations  were  carried  out  at  20°. 

PROCEDURE. 

Measurements  both  of  conductivity  and  viscosity  were  made  at  15°, 
25°,  and  35°.  The  data  were  calculated  in  the  usual  way,  tables  of 
constants  and  the  use  of  a  calculating  machine  greatly  facilitating  the 
operation.  The  viscosity  coefficients  were  obtained  from  the  formula 


where  ri0,  s0,  and  tQ  are  the  viscosity,  density,  and  time  of  flow  of  pure 
water,  and  T/,  s,  and  t  the  corresponding  values  for  the  liquid  in  question 
in  any  given  viscosimeter. 

Fluidity,  represented  by  <£,  is  equal  to  -•     The  temperature  coeffi- 

~n 

cients  in  conductivity  units  represent  the  actual  increase  in  molecular 
conductivity  per  degree  rise  in  temperature. 

Per  cent  temperature  coefficients,  both  of  conductivity  and  fluidity, 
were  calculated  from  the  formula : 


OF    CERTAIN    SALTS   IN   TERNARY   MIXTURES. 


123 


VISCOSITY   DATA. 

TABLE  33. — Density,   specific  conductivity,   viscosity    and  fluidity  of  ternary  mixtures   of 
glycerol,  acetone,  and  water  (G,  A,  W). 


Mix- 
ture? of 

15° 

25° 

35° 

G  A 

W 

d 

M 

. 

<£ 

d 

M 

f 

<*> 

d 

A* 

if 

0 

1  2 
1  2 

2 

.0063 
.0064 

0  .  00343 
.00215 

0.03539 
.03842 

28.26 
26.03 

0.9984 
0.9959 

0.00468 
.00294 

0.02530 
.02738 

39.53 
36.52 

0.9902 
0.9874 

0.00613 
.00390 

0.01888 
.02035 

52.87 
49.14 

1  1 

2 

.0443 

.00368  .04107)23.83 

1.0377 

.00508 

.02951133.89 

1.0308 

.00676 

.02174 

46.00 

1  1 

1 

.0515 

.00237 

.06379;15.68 

1.0439 

.00332 

.04349 

23.00 

1.0354 

.00453 

.04116 

32.09 

2  1 

2 

.0905 

.00233 

.0809012.36 

1.0839 

.00333 

.05414 

18.47 

1.0768 

.00458 

.03826 

26.14 

2  2 

1 

.0534 

.00139 

.08901 

11.24 

1.0453 

.00202 

.05954 

16.80 

1.0363 

.00283 

.05954 

23.94 

2  1 

1 

.1072 

.00132 

.  15296 

6.54 

1.0998 

.00202 

.09706 

10.31 

1.0923 

.00293 

.06474 

15.45 

TABLE  34. — Viscosity  and  fluidity  of  salts  in  the  1-2-2  solvent  (glycerol  1, 
acetone  2,  water  2),  at  15°,  25°,  36°. 

RUBIDIUM  BROMIDE  IN  THE  1-2-2  SOLVENT. 


Mol. 
cone. 

Viscosities. 

Fluidities. 

Temp,  coeff. 

1,15° 

7j25° 

i»35° 

015° 

025° 

035° 

15  to  25° 

25  to  35° 

0.60 
.25 
.10 
Solv. 

0.03571 
.03608 
.03506 
.03539 

0.02603 
.02555 
.02529 
.02530 

0.01950 
.01959 
.01887 
.01888 

28.00 
27.72 
28.52 
28.26 

38.42 
39.14 
39.54 
39.53 

51.28 
51.05 
52.99 

52.87 

0.0335 
.0412 
.0386 
.0399 

0.0371 
.0304 
.0340 
.0337 

RUBIDIUM  IODIDE  IN  THE  1-2-2  SOLVENT. 

0.50 

.25 
.10 
Solv. 

0.03477 
.03471 
.03551 
.03539 

0.02532 
.02494 
.02550 
.02530 

0.01921 
.01850 
.01888 
.01888 

28.76 
28.81 
28.16 
28.26 

39.49 
40.10 
39.22 
39.53 

52.06 
54.05 
52.97 

52.87 

0.0373 
.0390 
.0392 
.0399 

0.0318 
.0348 
.0350 
.0337 

AMMONIUM  IODIDE  IN  THE  1-2-2  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.03486 
.03483 
.03540 
.03539 

0.02552 
.02503 
.02548 
.02530 

0.01954 
.01860 
.01913 

.01888 

28.69 
28.71 
28.25 
28.26 

39.19 
39.95 
39.25 
39.53 

51.98 
53.74 
52.27 

52.87 

0.0366 
.0389 
.0389 
.0399 

0.0326 
.0345 
.0331 
.0337 

124 


CONDUCTIVITY   AND   VISCOSITY 


TABLE  35. — Viscosity  and  fluidity  of  salts  in  the  1-2-1  solvent  (glycerol  1, 
acetone  2,  water  1),  at  15°,  25°,  35°. 

RUBIDIUM  BROMIDE  IN  THE  1-2-1  SOLVENT. 


Mol. 
cone. 

-    Viscosities. 

Fluidities. 

Temp,  coeff. 

,15° 

,25° 

7;350 

015° 

025° 

#35° 

15  to  25e 

25  to  35° 

0.50 
.25 
.10 

Solv. 

0.04089 
.04004 
.03950 
.03842 

0.02950 
.02895 
.02837 
.02738 

0.02228 
.02180 
.02134 
.02035 

24.45 
24.98 
25.32 
26.03 

33.90 
34.54 
35.25 
36.52 

44.88 
45.87 
46.86 
49.14 

0.0387 
.0383 
.0392 
.0403 

0.0324 
.0325 
.0329 
.0345 

RUBIDIUM  IODIDE  IN  THE  1-2-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.03998 
.03960 
.03930 
.03842 

0.02878 
.02873 
.02819 
.02738 

0.02163 
.02162 
.02127 
.02035 

25.01 
25.25 
25.45 
26.03 

34.74 
34.81 
35.47 
36.52 

46.23 
46.25 
47.01 
49.14 

0.0389 
.0377 
.0394 
.0403 

0.0331 
.0300 
.0328 
.0345 

AMMONIUM  BROMIDE  IN  THE  1-2-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.04090 
.04064 
.03830 
.03842 

0.02941 
.02917 
.02843 
.02738 

0.02214 
.02196 
.02132 
.02035 

24.45 
24.61 
26.11 
26.03 

34.00 
34.28 
35.17 
36.52 

45.17 
45.54 
46.90 
49.14 

0.0391 
.0394 
.0347 
.0403 

0.0328 
.0328 
.0333 
.0345 

AMMONIUM  IODIDE  IN  THE  1-2-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.04032 
.03976 
.03928 
.03842 

0.02880 
.02864 
.02818 
.02738 

0.02166 
.  02154 
.02112 
.  02035 

24.80 
25.15 
25.46 
26.03 

34.72 
34.92 
35.49 
36.52 

46.17 
46.43 
47.35 
49.14 

0.0400 
.0388 
.0394 
.0403 

0.0330 
.0330 
.0334 
.0345 

TABLE  36. — Viscosity  and  fluidity  of  salts  in  the  1-1-2  solvent  (glycerol  1, 
acetone  1,  water  2),  at  15°,  25°,  36°, 

RUBIDIUM  BROMIDE  IN  THE  1-1-2  SOLVENT. 


Mol. 
cone. 

Viscosities. 

Fluidities. 

Temp,  coeff. 

,15° 

ij25° 

ij35° 

515° 

525° 

535° 

15  to  25° 

25  to  35° 

0.50 
.25 
.10 
Solv. 

0.04041 
.04131 
.04162 
.04107 

0.02857 
.02926 
.02940 
.02951 

0.02122 
.02175 
.02170 
.02174 

24.75 
24.21 
24.08 
23.83 

35.00 
34.18 
34.01 
33.89 

47.13 
45.98 
46.08 
46.00 

0.0414 
.0412 
.0412 
.0423 

0.0346 
.0345 
.0354 
.0358 

RUBIDIUM  IODIDE  IN  THE  1-1-2  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.03923 
.04068 
.04140 
.04107 

0.02781 
.02894 
.02905 
.02951 

0.02062 
.02152 
.02166 
.02174 

25.49 
24.58 
24.15 
23.83 

35.96 
34.55 
34.42 
33.89 

48.50 
46.47 
46.17 
46.00 

0.0416 
.0406 
.0425 
.0423 

0.0349 
.0345 
.0340 
.0358 

AMMONIUM  IODIDE  IN  THE  1-1-2  SOLVENT. 

0.50 

.25 
.10 
Solv. 

0.03926 
.04075 
.04148 
.04107 

0.02792 
.02885 
.02915 
.02951 

0.02060 
.02149 
.02153 
.02174 

25.47 
24.44 
24.11 
23.83 

35.82 
34.66 
34.31 
33.89 

48.54 
46.53 
46.23 
46.00 

0.0406 
.0418 
.0423 
.0423 

0.0355 
.0342 
.0348 
.0358 

AMMONIUM  BROMIDE  IN  THE  1-1-2  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.04050 
.04130 
.04170 
.04107 

0.02881 
.02933 
.02933 
.02951 

0.02126 
.02173 
.02176 
.02174 

24.69 
24.21 
23.98 
23.83 

34.71 
34.09 
34.09 
33.89 

47.04 
46.02 
45.96 
46.00 

0.0406 
.0408 
.0422 
.0423 

0.0355 
.0350 
.0348 
.0358 

OF    CERTAIN    SALTS   IN   TERNARY   MIXTURES. 


125 


TABLE  37.— Vi 

RUBIDIUM  BROMIDE  IN  THE  1-1-1  SOLVENT. 


of  salts  in  the  1-1-1  solvent  (glycerol  1, 
acetone  1,  water  1),  at  15°,  25°,  35°. 


Mol. 
cone. 

Viscosities. 

Fluidities. 

Temp,  coeff. 

1,15° 

1/25° 

7j350 

015° 

(£25° 

035' 

15  to  25° 

25  to  35° 

0.50 

.25 
.10 
Solv. 

0.06371 
.06358 
.06335 
.06379 

0.04404 
.04380 
.04309 
.04349 

0.03196 
.03158 
i     .03084 
i     .03116 

15.70 
15.73 
15.79 
15.68 

22.71 
22.83 
23.21 
23.00 

31.29 
31.67 
32.43 
32.09 

0.0447 
.0452 
.0470 
.0467 

0.0378 
.0387 
.0397 
.0395 

RUBIDIUM  IODIDE  IN  THE  1-1-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.06239 
.06234 
.06282 
.06379 

0.04402 
.04307 
.04340 
.04349 

0.03209 
.03085 
.03128 
.03116 

16.03 
16.04 
15.92 
15.68 

22.72 
23.22 
23.04 
23.00 

31.16 
32.42 
31.97 
32.09 

0.0417 
.0447 
.0447 
.0467 

0.0372 
.0396 
.0387 
.0395 

AMMONIUM  BROMIDE  IN  THE  1-1-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.06353 
.06338 
.06317 
.06379 

0.04397 
.04358 
.04304 
.04349 

0.03190 
.03143 
.03088 
.03116 

15.74 

15.78 
15.83 
15.68 

22.74 
22.05 
23.23 
23.00 

31.35 

31.82 
32.38 
32.09 

0.0445 
.0454 
.0468 
.0467 

0.0378 
.0385 
.0394 
.0395 

AMMONIUM  IODIDE  IN  THE  1-1-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.06119 
.06288 
.06338 
.06379 

0.04247 
.04333 
.04449 
.04349 

0.03054 
.03137 
.03153 
.03116 

16.34 
15.90 
15.78 
15.68 

23.55 
23.08 

22.48 
23.00 

32.74 
31.88 
31.72 
32.09 

0.0435 
.0451 
.0425 
.0467 

0.0391 
.0381 
.0411 
.0395 

TABLE  38. — Viscosity  and  fluidity  of  salts  in  the  2-1-2  solvent  (glycerol  2, 
acetone  1,  water  2),  at  15°,  25°,  35°. 

RUBIDIUM  BROMIDE  IN  THE  2-1-2  SOLVENT. 


Mol. 
cone. 

Viscosities. 

Fluidities. 

Temp,  coeff. 

,15° 

7725° 

,35° 

</>15° 

025° 

035° 

15  to  25° 

25  to  35° 

0.50 
.25 
.10 
Solv. 

0.07739 
.08014 
.08107 
.08090 

0.05334 
.05355 
.05432 
.05414 

0.03804 
.03782 
.03829 
.03826 

12.92 
12.48 
12.34 
12.36 

18.75 
18.67 
18.42 
18.47 

26.29 
26.44 
26.12 
26.14 

0.0450 
.0497 
.0493 
.0494 

0.0402 
.0416 
.0418 
.0415 

RUBIDIUM  IODIDE  IN  THE  2-1-2  SOLVENT. 

0.50 

.25 
.10 
Solv. 

0.07575 
.07730 
.07967 
.08090 

0.05165 
.05218 
.05370 
.05414 

0.03693 
.03654 
.03782 
.03826 

13.20 
12.94 
12.55 
12.36 

19.36 
19.16 
18.62 
18.47 

27.08 
27.37 
26.44 
26.14 

0.0467 
.0482 
.0484 
.0494 

0.0398 
.0427 
.0419 
.0415 

AMMONIUM  BROMIDE  IN  THE  2-1-2  SOLVENT. 

0.50 
j   .25 
<  .10 
ISolv. 

0.07739 
.07934 
.08054 
.08090 

0.05202 
.05366 
.05320 
.05414 

0.03702 
.03821 
.03763 
.03826 

12.92 
12.60 
12.42 
12.36 

19.22 
18.64 
18.80 
18.47 

27.01 
26.17 
26.57 
26.14 

0.0488 
.0479 
.0514 
.0494 

0.0405 
.0404 
.0418 
.0415 

AMMONIUM  IODIDE  IN  THE  2-1-2  SOLVENT. 

^0.50 
,  .25 

.10 
Solv. 

0.07545 
.07729 
.07854 
.08090 

0.05175 
.05206 
.05417 
.05414 

0.03683 
.03660 
.03855 
.03826 

13.25 
12.94 
12.73 
12.36 

19.32 
19.21 
18.46 
18.47 

27.15 
27.32 
25.94 
26.14 

0.0458 
.0484 
.0450 
.0494 

0.0405 
.0423 
.0405 
.0415 

126 


CONDUCTIVITY   AND    VISCOSITY 


TABLE  39. — Viscosity  and  fluidity  of  salts  in  the  2^2-1  solvent  (glycerol  2, 
acetone  2,  water  1),  at  15°,  25°,  35°. 

AMMONIUM  BROMIDE  IN  THE  2-2-1  SOLVENT. 


Mol. 
cone. 

Viscosities. 

Fluidities. 

Ternp.  coeff. 

,,15° 

1725° 

,35° 

015° 

#25° 

#35° 

15  to  25° 

25  to  35° 

0.50 
.25 
.10 
Solv. 

Gave  a  nc 
0.08989 
.08938 
.08901 

n-homoge 
0.06071 
.05906 
.  05954 

leous  solut 
0.04359 
.04178 
.04177 

ion.   Ac< 
11.12 
11.19 
11.24 

;tone  sal 
16.47 
16.94 
16.80 

ted  out. 
22.94 
23.83 
23.94 

0.0481 
.0514 
.0495 

0.0392 
.0407 
.0426 

AMMONIUM  IODIDE  IN  THE  2-2-1  SOLVENT., 

0.50 
.25 
.10 
Solv. 

0.08985 
.08939 
.08915 
.08901 

0.06493 
.06282 
.06290 
.05954 

0.04867 
.04437 
.04371 
.04177 

11.13 

11.19 
11.22 
11.24 

15.40 
15.92 
15.90 
16.80 

20.55 
22.54 
22.88 
23.94 

0.0384 
.0423 
.0418 
.0495 

0.0336 
.0416 
.0502 
.0426 

TABLE  40. — Viscosity  and  fluidity  of  salts  in  the  2-1-1  solvent  (glycerol  2, 
acetone  1,  water  1),  at  15°,  25°,  35°. 

RUBIDIUM  BROMIDE  IN  THE  2-1-1  SOLVENT. 


Mol. 
cone. 

Viscosities. 

Fluidities. 

Temp,  coeff. 

Tjl5° 

1,25° 

7?  35° 

<t>!5° 

025° 

#35° 

15  to  25° 

25  to  35° 

0.50 
.25 
.10 
Solv. 

0.15066 
.15388 
.  15462 
.  15296 

0.09583 
.09781 
.09812 
.09706 

0.06512 
.  06633 
.06623 
.06474 

6.64 
6.50 
6.47 
6.54 

10.44 
10.22 
10.19 
10.31 

15.36 
15.08 
15.10 
15.45 

0.0572 
.0573 
.0576 
.0576 

0.0491 
.0475 
.0481 
.0499 

RUBIDIUM  IODIDE  IN  THE  2-1-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.14625 
.15194 
.  15408 
.  15296 

0.09322 
.09682 
.09779 
.09706 

0.06316 
.06557 
.06589 
.06474 

6.84 
6.68 
6.49 
6.54 

10.73 
10.33 
10.23 
10.31 

15.83 
15.25 
15.18 
15.45 

0.0569 
.0546 
.0576 
.0576 

0.0476 
.0476 
.0482 
.0499 

AMMONIUM  BROMIDE  IN  THE  2-1-1  SOLVENT. 

0.50 

.25 
.10 
Solv. 

0.15177 
.  15417 
.  15550 
.  15296 

0  .  09650 
.09816 
.09814 
-.09706 

0.06575 
.06652 
.06624 
.06474 

6.63 
6.49 
6.45 
6.54 

10.36 
10.19 
10.19 
10.31 

15.21 
15.03 
15.10 
15.45 

0.0562 
.0571 
.0579 
.0576 

0.0468 
.0476 
.0482 
.0499 

AMMONIUM  IODIDE  IN  THE  2-1-1  SOLVENT. 

0.50 
.25 
.10 
Solv. 

0.14709 
.15211 
.  15410 
.  15296 

0.09366 
.09703 
.09757 
.09706 

0.06352 
.  06559 
.06588 
.06474 

6.80 
6.57 
6.49 
6.54 

10.68 
10.31 
10.25 
10.31 

15.74 
15.25 
15.08 
15.45 

0.0570 
.0568 
.0579 
.0576 

0.0475 
.0479 
.0471 
.0499 

OF    CERTAIN    SALTS   IN   TERNARY    MIXTURES. 


127 


CONDUCTIVITY   DATA. 

TABLE  41. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 
1-2-2  solvent  (glycerol  1,  acetone  2,  water  2),  at  15°,  25°,  35°. 

RUBIDIUM  BROMIDE  IN  THE  1-2-2  SOLVENT. 


Molecular 

Temperature  coefficients. 

V 

conductivities. 

Per  cent. 

Cond.  units. 

Mr  15° 

Mr  25° 

M«,35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 

23.50 

30.93 

39.20 

0.0316 

0.0267 

0.743 

0.827 

4 

23.14      29.33 

35.92 

.0267 

.0225 

.619 

.659 

10 

25.87      32.06 

39.39 

.0239 

.0242 

.619 

.659 

50 

29.10     37.09 

47.57 

.0275 

.0283 

.799 

1.048 

200 

30.60     39.84 

50.55        .0302 

.0269 

.924 

1.071 

RUBIDIUM  IODIDE  IN  THE  1-2-2  SOLVENT. 

2 

25.90  |  33.99 

42.76     0.0312 

0.0258     0.809 

0.877 

4 

25.74 

33.58 

43.85 

.0305 

.0306 

.784 

1.027 

10 

27.03     36.75 

.0360 

.972 

50 

30.08     40.08 

51.10 

.0332 

.0275 

1.000 

1.102 

200 

30.61  i  41.00 

52.49 

.0339 

.0280 

1.039 

1.149 

AMMONIUM  IODIDE  IN  THE  1—2—2  SOLVENT. 

2 

25.81 

33.87 

42.86 

0.0312 

0.0265 

0.806 

0.899 

4 

26.60 

34.99 

44.44 

.0315        .0270 

.839 

.945 

10 

28.72 

38.00 

48.70 

.0323 

.0282 

.928 

1.070 

50 

29.94 

39.78 

50.73 

.0329        .0275 

.984 

1.095 

200 

31.14 

41.69 

53.60 

.0339        .0286 

1  055 

1.101 

TABLE  42. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 
1-2-1  solvent  (glycerol  1,  acetone  2,  water  1),  at  15°,  25°,  36°. 

RUBIDIUM  BROMIDE  IN  THE  1-2-1  SOLVENT. 


V 

Molecular 
conductivities. 

Temperature  coefficients. 

Per  cent. 

Cond.  units. 

Mr  15° 

M,25° 

M,35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 
4 
10 
50 

14.78 
15.82 
17.24 
19.10 

19.35 
20.74 
22.69 
25.50 

24.49 
26.45 
29.12 
32.87 

0.0309 
.0311 
.0316 
.0335 

0.0266 
.0275 
.0288 
.0289 

0.457 
.492 
.545 
.640 

0.514 
.571 
.643 
.737 

RUBIDIUM  IODIDE  IN  THE  1-2-1  SOLVENT. 

2 
4 
10 
50 
200 
800 

18.33 
19.00 
20.00 
21.40 
22.56 
24.23 

23.81 
24.68 
26.22 
29.25 
29.85 
32.80 

30.24 
31.41 
33.74 
36.51 
38.69 
42.00 

0.0299 
.0299 
.0311 
.0322 
.0323 
.0354 

0.0270 
.0273 
.0287 
.0291 
.0296 
.0280 

0.548 
.568 
.622 
.688 
.729 
.857 

0.643 
.673 
.752 
.823 

.884 
.920 

AMMONIUM  BROMIDE  IN  THE  1-2-1  SOLVENT. 

2 
4 
10 
50 

15.94 
15.41 
17.34 
19.28 

19.72 
20.45 
22.81 

25.82 

24.94 
26.24 
29.23 
33.22 

0.0311 
.0327 
.0315 
.0339 

0.0265 
.0283 
.0282 
.0287 

0.468 
.504 
.547 
.654 

0.522 
.579 
.642 
.740 

AMMONIUM  IODIDE  IN  THE  1-2-1  SOLVENT. 

2 
4 

10 

18.23 
18.80 
19.88 

23.66 
24.67 
26.36 

30.28 
31.26 
33.54 

0.0298 
.0312 
.0326 

0.0283 
.0267 
.0272 

0.543 
.587 
.648 

0.662 
.659 

.718 

128 


CONDUCTIVITY   AND    VISCOSITY 


TABLE  43. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 
1-1-2  solvent  (glycerol  1,  acetone  1,  water  2),  at  15°,  25°,  35°. 

RUBIDIUM  BROMIDE  IN  THE  1—1—2  SOLVENT. 


Molecular 

Temperature  coefficients. 

V 

conductivities  . 

Per  cent. 

Cond.  units. 

Mr  15° 

M»25° 

M»35° 

15  to  25° 

25to35°15to25°;25to35° 

2 

26.35 

34.62 

44.31 

0.0314 

0.0280 

0.827 

0.969 

4 

25.59 

33.45 

41.99 

.0307 

.0255        .786 

.845 

10 

28.02 

37.23 

48.03 

.0329 

.0290        .921 

1.080 

50 

28.75 

39.86 

51.71 

.0386 

.0297      1.111 

1.185 

RUBIDIUM  IODIDE  IN  THE  1-1-2  SOLVENT. 

2 

27.42 

35.96 

45.94 

0.0311      0.0278     0.854 

0.998 

4 

26.80 

35.39 

45.36 

.0320  |      .0282  1     .859 

.997 

10 

28.17 

37.52 

48.67 

.0332        .0297  |     .935     1.115 

50 

29.17 

38.98 

51.16 

.0336 

.0312        .981      1.218 

800 

32.40 

43.92 

57.46 

.0356 

.0308     1.152  i  1.354 

AMMONIUM  BROMIDE  IN  THE  1-1-2  SOLVENT. 

2     26.57 

35.19 

44.71 

0.0324     0.0263 

0.862 

0.952 

4     25.98 

34.29 

43.03         .0320  i      .0255 

.831 

.874 

10 

27.85 

37.20 

47.92        .0336 

.0288 

.935 

1.072 

50 

29.40 

39.48 

51.17 

.0343 

.0296 

1.008 

1.169 

800     31.99 

43.36 

55.44        .0349 

.0279 

1.117 

1.208 

AMMONIUM  IODIDE  IN  THE  1-1-2  SOLVENT. 

2 

27.54 

35.86 

46.36 

0.0302  i  0.0293 

0.832 

1.050 

4     26.40 

35.22 

45.21 

.0334        .0284 

.882 

.999 

10     27.93 

37.62 

48.68 

.0347        .0294 

.969 

1.106 

50  i  28.88 

38.97 

50.67 

.0347        .0300 

1.001 

1.170 

200  !  30.82 

41.47 

54.02 

.0346        .0303 

1.065 

1.255 

800  i  31.86 

43.01 

56  .  34 

.0350        .0310 

1.115 

1.333 

TABLE  44. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 
1-1-1  solvent  (glycerol  1,  acetone  1,  water  1),  at  15°,  25°,  35°. 

RUBIDIUM  BROMIDE  IN  THE  1-1-1  SOLVENT. 


V 

Molecular 
conductivities. 

Temperature  coefficients. 

Per  cent. 

Cond.  units. 

M»15° 

M.25" 

Mr  35° 

15  to  25° 

25  to  35° 

15  to  25°  25  to  35° 

2 
4 
10 
50 
200 

15.22 

15.68 
16.51 
17.67 
18.19 

20.56 
21.20 
22.53 
24.41 
25.30 

26.80 
28.02 
29.71 
32.33 
33.43 

0.0351 
.0352 
.0365 
.0381 
.0391 

0.0304 
.0322 
.0319 
.0325 
.0321 

0  534 
552 
602 
674 
11 

0.624 
.682 
.718 
.794 
.813 

RUBIDIUM  IODIDE  IN  THE  1-1-1  SOLVENT. 

2 
4 
10 
50 
200 

16.07 
16.67 
17.65 

19!23 

21.95 
22.93 
24.24 

28.81 
29.98 
31.90 

0.0366 
.0375 
.0316 

0.0313 
.0307 
.0316 

0.588 
.626 
.659 

0.686 
.705 
.766 

.736 

26.73 

34.09 

.0389 

.0275 

.750 

OF    CERTAIN    SALTS   IN   TERNARY   MIXTURES. 


129 


TABLE  44. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 

1-1-1  solvent  (glycerol  1,  acetone  1,  water  1),  at  15°,  25°,  85°. — Continued. 

AMMONIUM  BROMIDE  IN  THE  1-1-1  SOLVENT. 


V 

Molecular 
conductivities. 

Temperature  coefficients. 

Per  cent. 

Cond.  units. 

Me  16° 

Mv25° 

^35° 

15  to  25° 

25  to  35° 

15to25° 

25  to  35° 

2 
4 
10 
50 
200 

15.21 
15.71 
16.56 
17.84 
18.50 

20.57 
21.41 
22.58 
24.59 
25.52 

27.11 

28.45 
30.03 
32.88 
34.13 

0.0352 
.0363 
.0363 
.0378 
.0379 

0.0418 
.0329 
.0326 
.0337 
.0338 

0.536 
.570 
.602 
.675 
.702 

0.654 
.704 
.745 
.829 
.863 

AMMONIUM  IODIDE  IN  THE  1-1-1  SOLVENT. 

2 
4 
10 
50 
200 

15.57 
16.86 
17.43 

17.87 
18.72 

20.51 
23.04 
23.68 
24.38 
26.05 

25.87 
30.35 
31.84 
33.20 
34.67 

0.0317 
.0366 
.0359 
.0364 
.0392 

0.0261 
.0317 
.0345 
.0362 
.0331 

0.494 
.618 
.625 
.651 
.733 

0.536 
.731 
.816 

.882 
.862 

TABLE  45. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 
2-1-2  solvent  (glycerol  2,  acetone  1,  water  2),  at  15°,  25°,  85°. 

RUBIDIUM  BROMIDE  IN  THE  2-1-2  SOLVENT. 


V 

Molecular 
conductivities. 

Temperature  coefficients. 

Per  cnet.              Cond.  units. 

M,15° 

M,25° 

M*35° 

15  to  25° 

25  to  35° 

15  to.  5°  25  to  35° 

2 
4 
10 
50 
200 

15.04 
15.55 
16.66 
17.54 
18.15 

20.75 
21.59 
22.99 
24.58 
25.54 

27.45 
28.66 
30.30 
32.92 
34.47 

0.0380 
.0388 
.0380 
.0401 
.0407 

0.0323 
.0327 
.0318 
.0339 
.0350 

0.571 
.604 
.633 
.704 
.739 

0.670 
.707 
.731 
.834 
.893 

RUBIDIUM  IODIDE  IN  THE  2-1-2  SOLVENT. 

2 
4 
10 
50 

200 

15.32 
15.57 
16.42 
17.07 
17.12 

21.29 
21.78 
22.98 
24.06 
24.24 

28.25 
29.02 
30.58 
32.38 
32.66 

0.0390 
.0399 
.0399 
.0409 
.0416 

0.0328  ;  0.597 
.0332        .621 
.0331        .656 
.0346        .699 
.0347  !     .712 

0.699 
.724 
.760 
.832 
.842 

AMMONIUM  BROMIDE  IN  THE  2-1-2  SOLVENT. 

2 
4 
10 
50 

200 

15.55 
15.60 
16.16 
16.92 
17.59 

21.59     28.55 
21.52  !  28.86 
22.52  I  30.13 
23.78     32.08 
24.73  !  33.25 

0.0388 
.0379 
.0394 
.0406 
.0406 

0.0322 
.0341 
.0338 
.0349 
.0345 

0.604 
.592 
.636 
.686 
.714 

0.696 
.734 
.761 
.830 
.852 

AMMONIUM  IODIDE  IN  THE  2-1-2  SOLVENT. 

2 
4 
10 
50 
200 

15.58 
15.80 
16.48 
17.11 
17.34 

21.71 
22.05 
22.61 
24.46 
24.57 

28.79 
29.39 
30.13 
32.88 
33.16 

0.0393 
.0396 
.0372 
.0430 
.0417 

0.0326 
.0333 
.0333 
.0344 
.0350 

0.613 
.625 
.613 
.735 
.723 

0.708 
.734 
.752 
.842 
.859 

130 


CONDUCTIVITY   AND   VISCOSITY 


TABLE  46. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 
2-2-1  solvent  (glycerol  2,  acetone  2,  water  1),  at  15°,  25°,  35°. 

AMMONIUM  BROMIDE  IN  THE  2-2-1  SOLVENT. 


v 

Molecular 
conductivities. 

Temperature  coefficients. 

Per  cent. 

Cond.  units. 

M*15° 

M,25° 

Me  35° 

15  to  25°J25  to  35° 

15to25°|25to35° 

2 
4 
10 
50 
200 

Gave  a 
9.09 
9.81 
10.62 
11.21 

non-hoi 
12.44 
13.71 
14.99 
15.93 

nogeneo 
16.85 
18.42 
20.43 
21.77 

us  solutio 
0.0369 
.0399 
.0411 
.0421 

a.     Aceto 
0.0355 
.0344 
.0363 
.0367 

ne  saltec 
0.335 
.391 
.437 
.472 

lout. 
0.441 
.471 
.544 
.584 

AMMONIUM  IODIDE  IN  THE  2—2—1  SOLVENT. 

2 
4 
10 
50 
200 

9.87 
10.63 
11.40 
12.18 
12.40 

13.23 
14.75 
15.80 
16.97 
17.50 

16.95 
19.68 
21.11 
23.12 
23.69 

0.0340 
.0388 
.0386 
.0393 
.0411 

0.0281 
.0334 
.0336 
.0362 
.0354 

0.336 
.412 
.440 
.479 
.510 

0.372 
.493 
.531 
.615 
.619 

TABLE  47. — Molecular  conductivities  and  temperature  coefficients  of  salts  in  the 
2-1-1  solvent  (glycerol  2,  acetone  1,  water  1),  at  15°,  25°,  35°. 

AMMONIUM  BROMIDE  IN  THE  2-1-1  SOLVENT. 


V 

Molecular 
conductivities. 

Temperature  coefficients. 

Per  cent. 

Cond.  units. 

Mr  15° 

M,25° 

Mr  35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 
4 
10 
50 
800 

7.91 
7.83 
8.23 
8.56 
9.19 

11.20 
11.25 
12.06 
12.67 
13.55 

15.82 
15.61 
16.88 
17.80 
19.34 

0.0416 
.0437 
.0465 
.0480 
.0474 

0.0413 
.0388 
.0400 
.0405 
.0427 

0.329 
.342 
.383 
.411 
.436 

0.462 
.436 
.482 
.513 
.579 

AMMONIUM  IODIDE  IN  THE  2-1-1  SOLVENT. 

2 
4 
10 
50 
800 

8.08 
8.16 

8.38 
8.87 
8.69 

11.79 
11.97 
12.29 
13.10 
12.84 

16.46 
16.77 
17.31 
18.55 
18.36 

0.0459 
.0467 
.0467 
.0477 
.0478 

0.0396 
.0401 
.0408 
.0416 
.0430 

0.371 
.381 
.391 
.423 
.415 

0.467 
.480 
.502 
.545 
.552 

RUBIDIUM  BROMIDE  IN  THE  2-1-1  SOLVENT. 

2 
4 
10 
50 
800 

7.62 
7.68 
8.10 
8.50 

11.19 
11.19 
11.93 
12.55 
12.62 

15.53 
15.57 
16.74 
17.72 
17.83 

0.0469 
.0453 
.0473 
.0476 

0.0388 
.0395 
.0403 
.0412 
.0413 

0.357 
.348 
.383 

.405 

0.434 
.441 
.481 
.517 
.521 

RUBIDIUM  IODIDE  IN  THE  2-1-1  SOLVENT. 

2 

4 
10 
50 
800 

7.92     11.31 
8.02  ;  11.79 
8.26  i  12.16 
8.53      12.67 
9.26  !  13.99 

16.18 
16.52 
17.10 
17.93 
20.01 

0.0420 
.0470 
.0472 
.0485 
.0511 

0.0433 
.0401 
.0406 
.0415 
.0430 

0.333 
.377 
.390 
.414 
.473 

0.487 
.473 
.494 
.526 
.602 

OF   CERTAIN   SALTS   IN  TERNARY   MIXTURES.  131 

DISCUSSION  OF  RESULTS. 

As  in  our  preceding  work  with  rubidium  and  ammonium  salts  in 
glycerol  and  acetone  and  their  binary  mixtures  with  water,  parallel 
investigations  in  viscosity  and  conductivity  have  been  carried  out. 

We  have  obtained  measurements  with  both  concentrated  and  moder- 
ately dilute  solutions  over  a  wide  range  of  ternary  mixtures  of  the  three 
solvents  employed.  Although  glycerol  and  acetone  of  themselves  are 
immiscible,  the  addition  of  about  20  per  cent  water  gives  a  perfectly 
homogeneous  liquid  which  corresponds  to  our  2-2-1  solvent.  However, 
it  should  be  noted  that  in  this  solvent  which  contains  the  lowest  per- 
centage of  water  in  the  series,  it  was  impossible  to  obtain  concentrated 
solutions  of  the  ammonium  salts  while  the  rubidium  salts  failed  to  go 
into  solution  at  concentrations  above  1/200  N.,  since  the  acetone 
immediately  separated  out  giving  a  non-homogeneous  solution. 

As  previously  noted,  table  33  shows  the  specific  data  relating  to  the 
various  mixtures,  viz,  the  density,  specific  conductivity,  viscosity,  and 
fluidity  at  all  three  temperatures  studied.  Reference  to  it  will  show 
that,  e.  g.,  at  25°,  the  standard  comparison  temperature,  the  solvents 
possess  densities  ranging  from  0.9984  for  the  1-2-2  mixture  to  1.0998 
for  the  2-1-1,  while  the  viscosities  for  the  same  solvents  vary  between 
0.02530  and  0.09706.  Thus,  the  viscosities  for  the  two  extremes  in 
the  series  lie  between  those  of  25  and  50  per  cent  glycerol  and  water 
(0.02017-0.06021)  for  the  former  and  between  the  75  and  50  per  cent 
(0.0135-0.06021)  for  the  latter  extreme.  The  values,  however,  lie  in 
both  cases  nearest  the  least  viscous  glycerol  mixture. 

The  specific  conductivity  of  the  1-1-1  solvent  is  about  3,000  times 
that  calculated  by  averaging  the  specific  conductivities  of  each  of  the 
constituents.  Jones  and  Davis1  have  noted  that  in  mixtures  of  glycerol 
and  water  containing  50  and  25  per  cent  glycerol,  the  specific  conduc- 
tivity was  higher  than  for  pure  water.  Their  explanation  is  that  the 
OH  ion  is  split  off;  i.  e.,  the  glycerol  is  dissociated  by  the  action  of  the 
water.  Jones  and  Bingham2  have  shown  that  the  molecular  conduc- 
tivity of  an  N/200  solution  of  potassium  iodide  in  acetone  is  about  the 
same  as  in  pure  water.  As  the  fluidity  of  acetone  is  about  2|  times  that 
of  water,  the  dissociating  action  of  acetone  would  be  of  the  order  of 
40  per  cent  that  of  water.  The  relative  association  factors  of  water 
and  acetone  would  lead  to  the  same  conclusion.  While  this  conclusion 
may  not  be  quantitatively  accurate,  it  is  safe  to  say  that_acetone  is  a 
strong  dissociating  agent.  It  is  therefore  possible  that  OH  ions  are 
split  off  from  the  glycerol  by  the  combined  action  of  the  water  and 
acetone,  and  possibly  some  from  the  water.  This  dissociation  would 
explain  the  very  high  specific  conductivities  of  the  solvents  used  in  this 
investigation  as  compared  with  those  calculated  by  averaging  the 
specific  conductivities  of  the  constituents.  

Carnegie  Inst.  WTash.  Pub.  No.  180.  "Ibid.,  No.  80. 


132  CONDUCTIVITY   AND   VISCOSITY 

Jones  and  Lindsay,1  continuing  the  investigation  of  the  phenomenon 
of  minima  in  conductivity  curves  observed  by  Zelinsky  and  Krapiwin 
and  by  Cohen,  advanced  the  theory,  based  on  the  hypothesis  of  Dutoit 
and  Aston,  that  the  decrease  in  conductivity  and  fluidity  in  solvents 
consisting  of  mixtures  of  associated  liquids,  was  due  to  the  fact  that 
each  liquid  decreased  the  association  of  the  other,  thus  decreasing  the 
size  of  the  ultimate  unit  particles  composing  the  solvent  and  increasing 
the  amount  of  frictional  surface  between  them.  With  these  two  con- 
ceptions in  mind;  i.  e.,  the  decrease  in  association  of  one  associated 
liquid  by  another  and  the  subsequent  effect  of  the  size  of  the  particles ; 
it  seems  clear  that  by  the  addition  of  a  third  associated  liquid  to  such 
a  binary  mixture,  the  decrease  in  association  would  be  carried  farther, 
resulting  in  an  increased  amount  of  frictional  surface  and  decreased 
fluidity.  If  the  unit  particles  in  the  acetone  were  much  larger  than  those 
already  in  the  binary  mixture,  the  fluidity  of  the  resulting  ternary  mix- 
ture would  be  increased.  This  is  not  probable,  as  acetone  is  an  asso- 
ciated liquid  and  the  presence  of  three  such  liquids,  each  decreasing 
the  association  of  the  others,  would,  it  is  reasonable  to  conclude,  result 
in  a  large  increase  in  the  number  of  smaller  particles.  All  conductivity 
and  fluidity  measurements  taken  during  this  investigation  support  this 
view.  The  comparisons  named  below  furnish  some  of  the  evidence  for 
these  conclusions. 

To  compare  the  results  obtained  with  those  calculated  from  averages, 
consider  the  data  obtained  by  Davis  and  Jones  with  glycerol-water 
mixtures  and  by  Davis,  Hughes,  and  Jones  with  acetone-water  mix- 
tures. They  used  rubidium  bromide  in  the  following  solvents:  75 
p.  ct.  glycerol  and  25  p.  ct.  water  (A) ;  75  p.  ct.  acetone  and  25  p.  ct. 
water  (B) ;  50  p.  ct.  glycerol  and  50  p.  ct.  water  (C) ;  50  p.  ct.  acetone 
and  50  p.  ct.  water  (D). 

It  should  be  noted  that  the  action  of  all  four  salts  used  in  this  inves- 
tigation do  not  differ  widely  in  the  same  solvents.  The  average 
obtained  with  the  2-1-1  and  1-2-1  solvents  can  be  compared  with  the 
A  and  B  solvents,  since  all  contain  25  per  cent  water.  The  75  per  cent  of 
ternary  mixture  is,  by  averaging  the  1-2-1  and  2-1-1  solvents,  equally 
divided  between  glycerol  and  acetone.  Solvents  C  and  D  can  be  com- 
pared with  the  1-1-2  solvents,  since  all  contain  50  per  cent  water. 

The  fluidity  of  the  1-1-1  solvent  is  about  one-sixth  that  calculated 
from  averages;  while  the  specific  conductivity  is,  as  noted  above,  about 
3,000  times  that  calculated  by  the  same  method;  hence  the  specific 
conductivity  of  the  solvent  is  18,000  times  that  which  would  be 
expected.  While  these  figures  can  be  considered  only  as  a  very  rough 
approximation,  they  indicate  a  relatively  large  dissociation  in  these 
ternary  mixtures. 

lCarnegie  Inst.  Wash.  Pub.  No.  80. 


OF    CERTAIN    SALTS   IN   TERNARY   MIXTURES.  133 

Further  evidence  for  this  view  is  obtained  from  the  1-2-2  solvent, 
whose  specific  conductivity  exceeds  by  a  much  larger  amount  that  cal- 
culated by  the  method  indicated  above.  The  relative  amounts  of 
acetone  and  water  are  much  larger  in  the  1-2-2  solvent  than  in  the 
1-1-1 ;  hence  a  larger  dissociation  of  the  glycerol  would  be  expected 
from  the  law  of  mass  action. 

The  viscosity  and  fluidity  tables  are  arranged  in  groups  under  each 
of  the  solvents.  Thus,  table  34  contains  the  data  for  rubidium  bromide, 
rubidium  iodide,  and  ammonium  iodide  in  the  1-2-2  solvent.  A  similar 
arrangement  is  carried  out  for  each  of  the  seven  solvents.  Associated 
with  each  table  of  viscosities  is  a  corresponding  table  of  temperature 
coefficients,  calculated  by  means  of  the  formula  given  on  page  122. 

It  has  been  shown  that  negative  viscosity  coefficients  occur  in  all 
cases  of  rubidium  salts  in  glycerol-water  mixtures,  and  also  for  ammo- 
nium bromide  and  iodide  in  these  solvents.  Such  was  also  found  to  be 
true  in  acetone-water  mixtures,  wherever  the  percentage  of  water  was 
higher  than  that  of  acetone. 

In  the  present  investigation  it  appears  that  a  similar  behavior  of 
such  salts  manifests  itself  wherever  the  solvents  are  of  the  same  general 
nature  as  those  mentioned  above.  Thus,  negative  viscosity  coefficients 
are  to  be  observed  in  the  case  of  all  the  salts  studied  in  the  1-1-1,  the 
2-1-2,  and  the  2-1-1  solvents.  Here  it  is  evident  that  either  glycerol 
or  water,  or  both,  are  present  in  greater  proportions  than  acetone. 
Since  glycerol  is  a  solvent  closely  allied  to  water  in  its  properties,  we 
may  disregard  its  enormous  viscosity  and  compare  these  solvents  with 
those  acetone-water  mixtures  in  which  the  water  is  present  in  the 
larger  proportion;  these  solvents  would  then  correspond  to  25,  20, 
and  25  per  cent  acetone-water  mixtures,  in  so  far  as  the  acetone  affects 
the  tendency  of  the  salts  to  lower  the  viscosity  of  the  solvent;  while  at 
the  same  time,  because  of  their  glycerol  content,  they  have  viscosity 
coefficients  comparable  with  50  to  25  per  cent  glycerol-water  mixtures. 

On  the  other  hand,  we  have  those  solvents  in  which  the  percentage 
content  of  acetone  exceeds  either  of  the  other  two  separately.  Under 
this  head  are  included  the  1-2-1  solvent  and,  in  certain  instances,  the 
solutions  in  the  2-2-1  mixture. 

In  the  case  of  the  1-2-2  and  the  1-1-1  solvents  an  apparent  fluctua- 
tion is  to  be  noted  in  the  concentration  curve  for  the  various  salts. 
Thus,  the  more  concentrated  solutions  increase  the  viscosity  of  the 
solvent,  while  the  more  dilute  lower  it;  a  possible  explanation  of  this 
phenomenon  is  suggested  later  in  discussing  the  conductivity  data. 
While  this  does  not  hold  for  all  temperatures,  it  appears  to  be  common 
to  all  the  salts  studied. 

Tables  41  to  47  show  the  molecular  conductivity,  temperature  coeffi- 
cients in  conductivity  units,  and  percentages  of  ammonium  iodide, 


134 


CONDUCTIVITY   AND   VISCOSITY 


ammonium  bromide,  rubidium  iodide,  and  rubidium  bromide  in  each 
of  the  solvents.  Figure  53  shows  the  conductivity  curves  of  ammonium 
iodide  in  the  ternary  solvents  at  25  degrees,  and  figure  54  shows  the 
corresponding  fluidity  curves.  While  these  curves  have  the  same 
general  character,  some  marked  differences  are  noticeable.  The  fluidi- 
ties of  glycerol,  water,  and  acetone  at  25  degrees  are  respectively  0.17, 
112.30,  and  288.95.  The  values  for  glycerol  and  water  are  taken  from 
the  data  of  Jones  and  Davis  and  those  for  acetone  from  the  work  of 
Jones  and  Bingham.  A  study  of  these  two  figures  in  the  light  of 
the  Thompson^Nernst2,  and  Dutoit- Aston3  hypotheses  will  afford  an 
explanation  of  all  cases  of  non-parallelism  in  the  two  sets  of  curves. 
Since  reducing  the  association  of  the  solvent  affects  both  its  fluidity 
and  dissociation,  and  since  the  relative  effect  on  each  is  not  known, 
the  above  explanation  is  to  be  regarded  as  only  qualitative. 


2-1-1  2-2-1  2-1-2  1-1-1  t-2-1  1-1-2  1-2-2 «- Solvents 

FIG.  53. — Conductivity  of  ammonium  iodide  in  glycerol,  acetone,  and  water,  at  25°. 

As  an  illustration,  consider  the  change  from  the  1-1-1  solvent  to  the 
1-2-1.  The  fluidity  increases  5  times  as  much  as  the  conductivity. 
The  changes  affecting  fluidity  are  as  follows :  the  water  changes  from 
33  to  25  per  cent  causing  a  small  reduction;  glycerol  changes  from 
33  to  25  per  cent  causing  a  rather  large  increase,  while  acetone  changes 
from  33  to  50  per  cent,  causing  a  very  large  increase.  The  changes  in 
water  and  glycerol  would  each  reduce  the  conductivity,  while  the 
change  in  acetone  would  increase  the  conductivity  much  less  than  the 
fluidity.  A  consideration  of  the  above  details  makes  it  clear  that  a 
much  larger  increase  in  fluidity  than  in  conductivity  should  be  expected. 

Other  changes  from  one  solvent  to  another  can  be  explained  by 
similar  considerations.  Thus,  figures  55  and  56  show  fluidity  and 
molecular  conductivity  curves  for  rubidium  bromide  in  the  1-2-1  and 
the  1-1-2  solvents  at  15°,  25°,  and  35°.  Jones,  Davis,  and  Hughes 
have  shown  that  for  glycerol-water  mixtures  and  acetone-water  mix- 
tures, temperature  coefficients  for  fluidity  are  larger  than  for  con- 
ductivity, because  rising  temperature  decreases  the  dissociation. 


l.  Mag.,  36,  320. 


2Zeit.  phys  Chem.,  13,  531.     (1894.) 


'Compt.  Rend.,  125,  240.     (1897.) 


OF   CERTAIN   SALTS   IN   TERNARY   MIXTURES.  135 

These  curves  show  that  the  same  relation  is  true  for  the  ternary  sol- 
vents. Sufficiently  dilute  solutions  have  not  been  used  in  this  work 
to  determine  the  dissociation  accurately,  no  measurements  beyond 
N/800  having  been  made.  The  data  obtained  by  Jones  and  Bingham, 
and  the  work  on  glycerol-water  and  acetone-water  mixtures  indicate 
that  solutions  more  dilute  than  N/1600  must  be  used.  Decrease  in 
dissociation  from  15°  to  35°  is  slight,  and  is  not  sufficient  to  explain 
the  difference  between  the  fluidity  and  conductivity  coefficients. 
Rubidium  and  ammonium  salts  are  not  in  the  class  of  salts  that  form 
complex  solvates,  yet  of  the  known  factors  which  affect  conductivity 
the  formation  of  solvates  is  the  only  one  which  can  explain  the  point 
here  raised.  If  a  solvate  is  formed  and  the  rise  in  temperature  reduces 
its  complexity  less  than  it  increases  the  fluidity  of  the  solvent,  the 
above  is  a  satisfactory  explanation.  In  this  connection  it  should  be 
noted  that  Jones  and  Guy1  have  found  some  evidence  for  the  formation 


2-1-1  2-2-1  2-1-2  1-1-1  1-2-1  1-1-2  1-2-2  «-  Solvent! 

FIG.  54. — Fluidity  of  smmonium  iodide  in  glycerol,  acetone,  and  water,  at  25°. 

of  glycerolates  by  sodium  and  potassium  salts.  The  acetone-water 
investigation  shows  some  evidence  for  the  formation  of  solvates  by 
rubidium  salts  in  a  mixed  solvent.  Another  factor  which  should  be 
investigated  is  the  polymerizing  action  of  acetone  and  the  effect  of 
temperature  on  the  complexity  of  the  polymers. 

It  should  be  noted  that  figures  55  and  56  show  two  distinct  types 
of  curves.  In  figure  55  both  conductivity  and  fluidity  curves  are 
regular,  while  in  figure  56  both  curves  show  a  minimum.  The  mini- 
mum occurs  at  the  N/4  point.  The  effect  on  conductivity  of  the  usual 
increase  in  dissociation  from  N/2  to  N/4  is  overcome  by  the  decrease 
in  fluidity,  thus  producing  the  minimum  point.  From  N/4  to  N/10 
there  is  little  change  in  fluidity,  hence  the  increase  in  dissociation  gives 
the  curves  a  sharp  upward  turn.  In  figure  55  the  influence  of  increas- 
ing fluidity  and  dissociation  work  together,  producing  a  convex  curve. 


iCarnegie  Inst.  Wash.  Pub.  No.  180. 


136 


CONDUCTIVITY   AND    VISCOSITY 


Returning  to  the  fundamental  point;  why  do  these  rubidium  and 
ammonium  salts  cause  conductivity  minima  and  also  fluidity  minima 
in  some  of  these  ternary  solvents  and  not  in  others?  In  some  cases, 
a  flat  curve  or  a  straight  line  is  produced.  In  all  such  cases  the  fluidity 
coefficients  are  negative.  These  negative  coefficients  occur  only  with 
the  solvents  containing  40  and  50  per  cent  acetone.  The  first  sug- 
gestion would  be  that  it  was  due  to  some  specific  effect  of  acetone  on 
the  fluidities  of  these  solvents,  but  a  study  of  them  and  also  of  figure 
58  renders  this  view  open  to  question.  The  explanation  suggested  is 
the  polymerizing  action  of  acetone.  The  normal  action  of  these  salts 
is  to  produce  positive  fluidity  coefficients  on  account  of  large  molecular 
volumes.  The  formation  of  a  polymer  would  reduce  the  number  of 


Volume  concentratii 


:  conductivity  curves.         =  fluidity  curves. 


FIG.  55. — Conductivity  and  fluidity  of  rubidium  bromide  in  the  1-2-1  solvent  at 
15°,  25°,  and  35°. 

molecules  and  hence  diminish  or  overcome  the  action  that  results  in 
these  positive  coefficients.  Jones  and  Mahin1  have  shown  that  cad- 
mium iodide,  lithium  nitrate,  and  lithium  acetate  polymerize  in  acetone. 
More  data  on  the  polymerization  of  inorganic  salts  is  desirable. 

Figure  57  shows  for  comparison  the  temperature  coefficients  of  con- 
ductivity and  fluidity  for  an  N/10  solution  of  ammonium  iodide  in  the 
ternary  solvents  at  15°  to  25°  and  25°  to  35°.  There  is  a  striking 
similarity  between  the  curves  for  the  conductivity  and  fluidity  coef- 
ficients.2 They  are  lower  for  the  higher  range  in  temperature,  as 

'Carnegie  Inst.  Wash.  Pub.  No.  180. 

2The  unusual  feature  shown  by  curve  II  for  the  2-2-1  solvent  is  probably  due  to  this  solvent 
containing  40  per  cent  of  acetone,  a  very  volatile  liquid.  This  solvent  also  contains  only  20  per 
cent  of  water,  which  is  the  minimum  amount  necessary  to  cause  these  three  liquids  to  form  a 
homogeneous  mixture. 


OF    CERTAIN    SALTS   IN   TERNARY   MIXTURES. 


137 


would  be  expected  from  the  work  in  other  solvents.  The  temperature 
coefficients  of  conductivity  are  very  close  to  those  calculated  from 
averages  from  the  data  of  the  glycerol-water  and  acetone-water  investi- 
gations. As  glycerol  has  a  much  higher  temperature  coefficient  of 
conductivity  than  either  acetone  or  water,  the  solvents  containing 
the  largest  percentage  of  glycerol  should  have  the  highest  coefficients. 
The  curves  are  in  keeping  with  this  fact. 

Conductivity  and  fluidity  in  these  ternary  solvents  is  much  below 
the  average.  These  considerations  emphasize  the  fact  that  fluidity 
outweighs  all  the  other  factors  that  affect  conductivity.  Furthermore, 
as  already  noted,  the  fluidity  data  indicate  that  association  is  more 


10 

=  conductivity  curves. 


:  fluidity  curves. 


FIG.  56. — Conductivity  and  fluidity  of  rubidium  bromide  in  the  1-1-2  solvent  at 
5°,  25°,  and  35°. 

reduced  in  the  ternary  than  in  the  binary  solvents,  but  a  decreased 
dissociation  is  not  indicated,  as  might  be  expected  from  the  deduction 
of  Jones  and  Lindsay. 

It  is  important  to  determine  what  effect  glycerol,  acetone,  and  water 
have  on  each  other  when  constituting  a  ternary  mixture.  This  investi- 
gation has  shown  that  the  properties  of  these  ternary  solvents  are 
widely  different  from  those  which  can  be  calculated  from  averages. 
The  curves  of  figure  58  are  drawn  to  show  the  differences  between  the 
measured  and  calculated  conductivities  and  fluidities.  To  illustrate: 
the  fluidity  of  the  2-1-1  solvent  is  10.31  at  25°;  calculated  from  averages 
it  is  100.4;  the  measured  is  thus  10  per  cent  of  the  calculated;  hence 
10  is  the  ordinate  for  the  2-1-1  point.  The  data  for  drawing  the  con- 


138 


CONDUCTIVITY   AND   VISCOSITY 


ductivity  curves  are  not  as  full  as  could  be  desired.  For  acetone  the 
value  is  taken  from  the  data  of  Jones  and  Bingham1  on  potassium 
iodide  for  a  N/200  solution  at  25°.  The  data  for  rubidium  bromide 
in  glycerol  and  in  water  are  taken  from  the  work  of  Jones  and  Davis.2 
The  action  of  potassium  and  rubidium  salts  in  relation  to  conductivity 
are  similar  enough  for  the  purpose  of  this  comparison.  The  2-1-1 
ordinate  shows  that  the  measured  conductivity  is  19.2  per  cent  of  the 
calculated.  That  the  curves  have  the  same  character  is  another  evi- 
dence for  the  close  relation  existing  between  conductivity  and  fluidity. 
In  figure  58  the  solvents  are  arranged  from  left  to  right  in  the  order 
of  the  percentage  of  glycerol  which  they  contain;  hence  the  curves 
show  that  the  more  glycerol  the  solvents  contain,  the  more  the  con- 


Conductivity  and  fluidity  coefficients  multiplied  hy  100 

t_3iS_8_8_8__S_8_|_*3_J 

1  F 

II  F 
III  C 

rvc 

uidity  coefficients  15'  -25" 
uidity  coefficients  25c-35° 
anduct  vity  coefficients  15'-25° 
>nductivity  coefficients  25'-35'  - 

olvenU 

\ 

\ 

z 

\ 

3 

\ 

n 

2 

^ 

\ 

•2~ 

\\ 

/ 

\ 

•'" 

\ 

IV 

•:A 

^^s 

Ss^ 

\ 

\ 

,.' 

2-1-1    2-2-1    2-1-2    1-1-1    I-! 

A    1-1-2    1-5 

-2  -  £ 

FIG.  57. — Conductivity  and  fluidity  temperature  coefficients 
for  N/10  solution  of  ammonium  iodide  in  glycerol, 
acetone,  and  water. 

ductivity  and  fluidity  values  depart  from  the  calculatedTaverages. 
The  theory  of  Jones  and  Veazey  states  that  viscosity  is  due  to  the 
friction  between  the  particles  of  the  liquid.  It  is  clear  that  the  smaller 
the  particles  the  greater  will  be  the  amount  of  frictional  j  surf  ace 
between  them;  hence,  the  greater  the  viscosity  of  any  homogeneous 
liquid,  the  smaller  must  be  the  particles  composing  it.  The  density 
of  the  liquid  should  also  affect  the  viscosity.  The  densities  of  glycerol, 
water,  and  acetone  are  1.26,  1.00  and  0.79,  respectively,  while  the 
fluidities  are  in  the  ratio  1  : 702 : 1741 ;  hence  the  variation  in  density 
is  so  small  in  comparison  with  the  variation  in  fluidity  that  the  former 
can  be  neglected.  It  seems  probable  that  there  is  one  other  important 
factor  that  affects  fluidity,  which  must  be  considered  in  addition  to 
the  size  of  the  particle.  The  particle  in  a  pure  homogeneous  liquid 
would  be  either  one  molecule  or  an  association  of  molecules. 


Carnegie  Inst.  Wash.  Pub.  No.  80. 


*IUd.,  No.  180. 


OF    CERTAIN    SALTS   IN   TERNARY   MIXTURES. 


139 


The  conception  of  molecular  volume  is  opposed  to  the  view  that 
the  glycerol  particle  is  smaller  than  that  of  acetone  or  water.  Acetone 
and  glycerol  have  the  same  molecular  volume,  calculated  on  the  basis 
of  the  simple  molecule,  i.  e.,73;  water  has  18.  Considering  the  associa- 
tion factors,  the  molecular  values  are  as  follows:  glycerol  150,  acetone 
92,  water  72.  From  the  latter  consideration,  glycerol  has  the  larger 
molecular  volume,  and  hence  should  have  the  highest  fluidity,  which 
is  contrary  to  the  facts.  The  error  probably  arises  from  the  use  of 
the  density  factor  in  calculating  molecular  volume.  The  density  of  a 
liquid  is  affected  by  both  the  density  of  the  molecules  and  the  spaces 
between  them.  The  kinetic  molecular  hypothesis  states  that  for  a  gas 
the  space  between  the  molecules  is  much  greater  than  that  occupied 
by  the  molecules.  For  a  liquid,  the  intermolecular  space  is  less  than 
for  a  gas,  and  for  a  solid  less  than  for  a  liquid.  For  liquids  it  seems 


Solvents—*  2-1-1  2-2-1  2-1-2 

FIG.  58. — Ordinates  are  percentages  by  which  measured  conductivity  and  fluidity 

differ  from  the  values  calculated  by  averages. 
Curve  I,  calculated  from  conductivity  of  Rbl  at  25°. 
Curve  II,  calculated  from  fluidity  of  solvents  at  25°. 

reasonable  to  believe  that  the  density  of  the  liquid  gives  no  idea  of  the 
density  of  the  molecule  or  the  space  between  them;  and  it  is  the  latter 
factor  of  intermolecular  space  as  well  as  the  size  of  the  particle  that 
must  be  considered  as  the  chief  factor  governing  fluidity.  Hence  the 
conception  of  molecular  volume  as  applied  to  liquids  is  not  opposed 
to  the  theory  that  fluidity  depends  on  the  frictional  surface  between 
the  particles  or  molecular  associations,  and  the  necessary  corollary 
that  the  amount  of  frictional  surface  depends  on  the  size  of  the  par- 
ticles. It  should  be  noted  in  this  connection  that  it  is  clear  that  the 
intermolecular  space  of  a  gas  is  the  most  important  factor  affecting 
its  viscosity.  It  is  reasonable  to  believe  that  the  same  factor  must 
be  considered  in  studying  the  viscosity  or  fluidity  of  liquids,  although 
it  is  relatively  less  important  than  for  gases.  It  is  probable  that  the 
present  conception  of  molecular  volume  as  applied  to  solids  is  approxi- 
mately correct,  since  the  intermolecular  space  may  be  so  small  that  it 
can  be  neglected. 


140          CONDUCTIVITY    AND    VISCOSITY    OF    CERTAIN    SALTS. 

SUMMARY. 

1.  The  conductivities  of  the  ternary  solvents  make  it  probable  that 
water  and  acetone  act  as  dissociating  agents  on  glycerol. 

2.  The  decrease  in  dissociation  of  one  associated  liquid  by  another  is 
much  larger  in  a  ternary  than  in  a  binary  mixture,  thus  producing 
decreased  values  in  conductivity  and  fluidity. 

3.  A  consideration  of  the  hypotheses  of  Dutoit  and  Aston  and  of 
Thompson-Nernst,  together  with  the  fluidities  of  glycerol,  acetone,  and 
water,  explain  the  differences  between  the  conductivity  and  fluidity 
curves  in  these  ternary  solvents. 

4.  Temperature  coefficients  of  fluidity  are  larger  than  for  conduc- 
tivity, as  in  binary  solvents.     The  formation  of  solvates  is  a  possible 
explanation  of  the  difference  between  these  coefficients. 

5.  The  minimum  point  which  occurs  in  some  of  the  conductivity 
curves  is  explained  by  considering  the  fluidity  of  the  solution. 

6.  A  possible  explanation  of  the  fluidity  changes  which  produce 
minima  in  the  conductivity  and  fluidity  curves  is  the  polymerization 
of  the  salts  by  the  acetone. 

7.  The  conductivity  and  fluidity  values  of  the  solvents  containing 
the  largest  percentage  of  glycerol  are  farthest  below  the  values  calcu- 
lated by  averages. 

8.  The  temperature  coefficients  of  conductivity  are  about  the  same 
as  those  calculated  from  averages. 

9.  This   investigation   emphasizes   the   fact   already   known   that 
fluidity  probably  outweighs  all  the  other  factors  affecting  conductivity. 

10.  The  conditions  governing  viscosity  in  a  pure  homogeneous  liquid 
are  discussed  from  a  physical  point  of  view. 


CHAPTER  VII. 

DISCUSSION  OF  EVIDENCE  ON  THE  SOLVATE  THEORY  OF  SOLUTION 

OBTAINED  IN  THE  LABORATORIES  OF  THE  JOHNS 

HOPKINS  UNIVERSITY.1 

About  fifteen  years  ago  the  work  which  led  to  the  present  theory  of 
solution  was  begun  in  this  laboratory.  From  a  very  simple  begin- 
ning, which  did  not  have  for  its  object  the  study  of  the  nature  of 
solution  in  general,  the  work  has  widened  in  a  fairly  large  number  of 
directions.  There  have  already  been  published  by  my  co-workers 
and  myself  about  eighty  papers  dealing  with  one  or  another  phase 
of  the  problem.  These  are  fairly  widely  scattered  through  chemical 
and  physical  literature,  having  been  published  in  American,  German, 
French,  and  English  scientific  journals.  In  addition,  the  Carnegie 
Institution  of  Washington,  which  has  so  generously  supported  the 
work,  and  without  which  it  would  have  been  impossible  to  have  car- 
ried out  many  of  the  investigations,  has  published  nine  monographs 
of  researches  bearing  directly  and  indirectly  on  the  question  of  the 
nature  of  solution. 

Taking  all  of  these  facts  into  account,  it  has  seemed  desirable  to 
discuss  here,  as  briefly  as  possible,  the  more  important  lines  of  evidence 
which  have  been  brought  out,  bearing  on  the  nature  of  that  condition 
of  matter  which  gives  rise  to  the  sciences  of  chemistry,  geology,  and 
to  a  large  part  of  biology. 

EARLIER  WORK. 

In  the  summer  of  1893  I  went  to  Stockholm  to  work  with  Svante 
Arrhenius.  He  suggested  that  we  carry  out  a  research  on  the  question 
as  to  whether  sulphuric  acid  forms  a  few  definite  hydrates  when  in  the 
presence  of  water,  as  the  theory  of  Mendele"eff  maintained.  According 
to  this  theory,  some  of  these  hydrates  were  very  complex,  one  of 
them  containing  as  much  as  100  molecules  of  water  to  1  molecule  of 
sulphuric  acid.  Mendeleeff  arrived  at  this  conclusion  chiefly  from  a 
study  of  the  specific  gravities  of  aqueous  solutions  of  sulphuric  acid  of 
different  concentrations.  The  specific-gravity  curves  showed  certain 
discontinuities  or  breaks,  which  Mendeleeff  interpreted  to  mean  the 
existence  of  definite  chemical  compounds  or  hydrates. 

At  the  suggestion  of  Arrhenius,  I  studied  the  problem  in  the  following 
way.  Acetic  acid  was  used  as  the  solvent.  The  freezing-point  lowerings 
of  the  acid  produced  by  adding  different  known  amounts  of  water  were 
determined.  The  freezing-point  lowerings  produced  by  adding  known 
amounts  of  sulphuric  acid  to  pure  acetic  acid  were  measured.  The 
freezing-point  lowerings  produced  by  adding,  simultaneously,  known 
amounts  of  sulphuric  acid  and  known  amounts  of  water  to  known 

'See  paper  in  Journ.  Franklin  Institute,  Nov.  and  Dec.  1913. 

141 


142  DISCUSSION   OF   EVIDENCE. 

amounts  of  acetic  acid  were  also  determined.  By  comparing  the  three 
sets  of  results,  from  Raoult's  law  we  could  calculate  the  composition 
of  the  hydrates  formed  by  the  sulphuric  acid,  if  any  were  formed. 

We  added  large  amounts  of  water  relative  to  the  sulphuric  acid, 
but  could  obtain  no  evidence  of  any  hydrate  of  sulphuric  acid  more 
complex  than  H2SO4,2H2O,  which  is  the  well-known  compound  H6SO6. 
We  did  not  obtain  the  slightest  evidence  of  the  existence  of  any  of  the 
more  complex  hydrates  which  Mendeleeff,  from  his  work,  had  supposed 
to  exist.  I  was,  therefore,  at  that  time,  thoroughly  convinced  of  the 
untenability  of  the  Mendele"eff  theory  of  hydrates,  and,  indeed,  of  any 
theory  of  hydration.  I  regarded  the  ions  in  solution  as  having  an 
existence  not  only  independent  of  one  another,  but  also  independent 
of  the  molecules  of  the  solvent.  In  a  word,  I  was  at  that  time  firmly 
convinced  that  no  theory  of  hydration  was  necessaiy  to  explain  the 
facts  that  were  then  known.  This  seemed  to  be  the  view  which  was 
held  at  that  time  also  by  most  of  those  who  founded  the  new  school 
of  chemistry. 

A  comparatively  few  years  later,  my  cooperators,  Ota1  and  Knight,2 
brought  to  light  certain  facts  which  could  not  be  explained  in  terms  of 
any  relation  that  was  then  known.  They  found  that  certain  double 
salts,  such  as  double  chlorides,  nitrates,  sulphates,  cyanides,  etc.,  pro- 
duced abnormally  great  lowering  of  the  freezing-point  of  water  when 
the  solutions  were  concentrated.  What  was  more  perplexing  was  the 
fact  that  the  molecular  depression  of  the  freezing-point  increased  with 
the  concentration  beyond  a  certain  definite  concentration. 

Similar  results  were  found  for  a  fairly  large  number  of  salts  by  Jones 
and  Chambers,3  and  by  Chambers  and  Frazer  working  with  Jones.4 
The  salts  studied  by  these  workers  were  those  that  are  known  to  be 
very  hydroscopic,  to  have  great  power  of  combining  with  water.  The 
question  arose,  what  did  these  results  mean?  At  that  time  I  was 
antagonistic  to  any  hydrate  theory.  My  experience  in  the  laboratory 
of  Arrhenius  had  produced  that  state  of  mind;  yet  I  was  unable  to 
explain  our  results  in  terms  of  any  other  assumption  than  that  a  part 
of  the  water  present  was  combined  with  the  dissolved  substance,  and 
was  therefore  removed  from  playing  the  role  of  solvent.  I  ventured 
this  suggestion,  for  want  of  any  better  in  1900.5 

The  suggestion  of  hydration  in  aqueous  solution  would  explain  the 
results  that  had  been  obtained.  If  a  part  of  the  water  present  was 
combined  with  the  dissolved  substance,  there  would  be  less  water  act- 
ing as  solvent ;  and  since  freezing-point  lowering  is  proportional  to  the 
ratio  between  the  number  of  molecules  of  the  solvent  and  of  the  dis- 
solved substance,  the  less  solvent  present  the  greater  the  lowering  of 
its  freezing-point.  It  is  one  thing  to  make  a  suggestion  which  accounts 
for  the  known  facts;  it  is  a  very  different  matter  to  show  that  this  is 

1Amer.  Chem.  Journ.,  22,  5  (1899).  3Ibid.,  23,  89  (1900).  6Ibid.,  23,  103  (1900). 

*lUd.,  110  (1899).  *Ibid.,  512  (1900). 


DISCUSSION   OF   EVIDENCE.  143 

the  only  reasonable  suggestion  which  will  account  for  them,  to  show 
that  the  suggestion  is  true. 

Aided  by  a  grant  from  the  Carnegie  Institution  of  Washington,  I 
started  Dr.  Getman1  on  a  more  or  less  systematic  study  of  the  whole 
problem.  The  question  arose,  were  the  results  already  obtained 
limited  to  a  few  compounds,  or  types  of  compounds,  or  was  this  a 
general  phenomenon?  We  took  up  the  study  of  acids,  bases,  and  salts 
in  concentrated  solutions,  especially  by  the  freezing-point  and  conduc- 
tivity methods.  We  also  studied  the  refractivities  of  many  solutions. 

RELATION  BETWEEN  LOWERING  OF  THE  FREEZING-POINT  OF  WATER  AND  WATER 
OF  CRYSTALLIZATION  OF  THE  DISSOLVED  SUBSTANCE. 

The  work  of  Getman  included  the  study  of  the  lowering  of  the  freez- 
ing-point of  water  produced  by  concentrated  solutions  of  the  chlorides 
of  sodium,  potassium,  ammonium,  lithium,  barium,  strontium,  cal- 
cium, magnesium,  iron,  and  aluminium;  the  bromides  of  sodium, 
potassium,  lithium,  barium,  strontium,  calcium,  magnesium,  and  cad- 
mium; the  iodides  of  sodium,  potassium,  lithium,  calcium,  barium, 
strontium,  and  cadmium;  and  the  nitrates  of  sodium,  potassium, 
ammonium,  lithium,  calcium,  magnesium,  manganese,  cobalt,  nickel, 
cadmium,  zinc,  aluminium,  iron,  and  chromium.  The  relation  between 
lowering  of  freezing-point  and  water  of  crystallization  can  be  seen  very- 
well  from  the  curves.2 

The  nitrates  of  sodium,  potassium,  and  ammonium,  which  crystallize 
without  water,  produce  the  smallest  lowering  of  the  freezing-point  of 
water.  Then  come  the  nitrate  of  lithium  with  2  molecules  of  water, 
calcium  with  4,  and  a  large  number  of  nitrates  each  with  6  molecules 
of  crystal  water;  all  give  about  the  same  lowering  of  the  freezing-point. 
Finally,  the  three  nitrates  of  aluminium,  iron,  and  chromium  with  8 
and  9  molecules  of  water,  give  the  greatest  lowering  of  the  freezing- 
point  of  water. 

Relations  similar  to  the  above  come  out  for  the  chlorides,  the 
bromides,  and  the  iodides.3  The  freezing-point  lowerings  of  water 
produced  by  them  are  roughly  proportional  to  the  amounts  of  water 
with  which  the  salts  crystallize. 

If,  on  the  other  hand,  we  compare  the  chlorides  with  the  bromides, 
with  the  iodides,  with  the  nitrates,  similar  relations  manifest  themselves. 

It  was  found  that  chlorides,  bromides,  iodides,  and  nitrates  which 
have  no  water  of  crystallization,  all  produce  about  the  same  molecular 
lowering  of  the  freezing-point  of  water,  and  this  is  between  3  and  4. 

lAmer.  Chem.  Journ.,  27,  433  (1902) ;  31,  303  (1904) ;  32,  308  (1904) ;  Zeit.  phys.  Chem..  46,  244 
(1903);  49,  385  (1904);  Phys.  Rev.,  18,  146  (1904);  Ber.  d.  chem.  Gesell.,  37,  1511  (1904). 
2See  Carnegie  Inst.  Wash.  Pub.  No.  60,  p.  24. 
3See  Carnegie  Inst.  Wash.  Pub.  No.  60,  pp.  20-26. 


144  DISCUSSION    OF   EVIDENCE. 

With  salts  that  crystallize  without  water  there  is  only  a  very  slight 
increase  in  the  molecular  lowering  of  the  freezing-point  with  increase 
in  the  concentration  of  the  solution.  The  salts  of  lithium,  which 
crystallize  with  the  same  amounts  of  water,  give  approximately  the 
same  depressions  of  the  freezing-point. 

If  we  compare  the  salts  of  the  alkaline  earths  that  crystallize  with 
6  molecules  of  water,  they  produce  approximately  the  same  lowerings ; 
the  nitrates  of  iron  and  aluminium  with  8  and  9  molecules  of  water 
give  greater  lowerings  than  the  corresponding  halogens  with  6. 

In  the  first  case  we  have  kept  the  acid  constant  and  compared  with 
one  another  the  salts  of  the  different  metals  with  the  same  acid.  In  the 
second  case  we  have  kept  the  metal  constant,  and  compared  the  salts  of 
a  given  metal  with  different  acids.  In  both  cases  the  relation  between 
lowering  of  the  freezing-point  of  water  by  the  dissolved  substance  and 
water  of  crystallization  of  the  dissolved  substance  manifests  itself. 

These  salts  that  crystallize  with  the  largest  amounts  of  water  produce 
the  greatest  molecular  lowering  of  the  freezing-point  of  water.  The 
work  was  done  with  concentrated  solutions,  and  it  has  already  been 
pointed  out  that  for  such  substances  the  molecular  lowering  of  the 
freezing-point  increases  with  the  concentration  of  the  solution. 

We  must  now  ask  what  bearing  has  this  relation  on  the  question  of 
hydration  or  non-hydration  in  aqueous  solution?  A  moment's  thought 
will  show  that  the  bearing  is  a  very  direct  one.  If  hydrates  exist  in 
aqueous  solution,  those  substances  which  in  such  solutions  would  form 
the  most  complex  hydrates  would  be  the  substances  that  would  crystal- 
lize from  aqueous  solutions  with  the  largest  amounts  of  water.  This  is 
the  same  as  to  say  that  those  substances  which,  in  the  presence  of  a  large 
amount  of  water,  have  the  greatest  power  to  combine  with  water,  would, 
other  things  being  equal,  be  the  ones  to  bring  with  them,  out  of  aqueous 
solution,  the  largest  amounts  of  water  as  water  of  crystallization. 

We  could  not,  however,  expect  one  of  these  phenomena  to  be  strictly 
a  linear  function  of  the  other,  since  there  are  undoubtedly  other  factors, 
such  as  shape  of  molecules,  angles  of  crystals,  etc.,  coming  into  play 
in  determining  the  exact  composition  of  crystals. 

That  a  relation  such  as  was  pointed  out  above  holds  so  well  and  so 
generally  for  such  a  large  number  of  substances  is  very  significant 
and  early  led  me  to  believe  that  the  suggestion  of  hydration  in  general 
in  aqueous  solution  contained  more  truth  than  I  imagined  when  it 
was  first  suggested. 

Having  found  a  relation  such  as  the  above,  we  were  led  to  look  about 
for  others  that  would  bear  directly  or  indirectly  on  the  problem  in 
hand.  Before  taking  up  these,  another  feature  of  the  work  of  Getman 
must  be  briefly  discussed. 


DISCUSSION   OF   EVIDENCE.  145 

APPROXIMATE  COMPOSITION  OF  THE  HYDRATES  FORMED  BY  VARIOUS 
SUBSTANCES  IN  SOLUTION. 

The  line  of  evidence  just  discussed  seemed  so  strongly  in  favor  of 
the  general  correctness  of  the  view  that  there  is  combination  between 
the  dissolved  substance  and  some  of  the  water  present,  that  Jones 
and  Getman1  undertook  to  calculate  the  approximate  composition  of 
the  hydrates  formed  by  the  different  substances,  and  by  the  same 
substance  at  different  dilutions. 

The  experimental  work  consisted  in  determining  the  freezing-point 
of  the  solution  and,  consequently,  the  depression  of  the  freezing-point 
of  water  produced  by  the  dissolved  substance  at  the  concentration  in 
question.  From  the  freezing-point  lowering  the  molecular  lowering 
was  calculated. 

The  dissociation  of  the  solution  was  measured  by  means  of  the 
conductivity  method.  Knowing  the  dissociation,  the  theoretical 
molecular  lowering  was  calculated  on  the  assumption  that  none  of 
the  solvent  was  combined  with  the  dissolved  substance.  The  ratio 
of  the  theoretical  molecular  lowering  to  the  value  found  experimentally, 
gave  the  proportion  of  all  the  water  present  that  was  uncombined. 
The  remainder  of  the  water  was,  of  course,  combined  with  the  dissolved 
substance.  The  total  amount  of  water  present  in  any  given  solu- 
tion could  be  readily  determined.  It  was  only  necessary  to  take 
the  specific  gravity  of  the  solution  by  weighing  a  known  volume  of  it. 
Knowing  the  specific  gravity  and  the  concentration,  it  was,  of  course, 
perfectly  simple  to  determine  the  total  amount  of  water  in,  say,  a 
liter  of  the  solution.  The  total  amount  of  water  in  the  solution  and 
the  percentage  of  combined  water  being  known,  the  total  amount 
of  combined  water  was  known.  Knowing  the  amount  of  dissolved 
substance  present  in,  say,  a  liter  of  the  solution,  and  knowing  the  total 
amount  of  water  combined  with  it,  it  was  perfectly  simple,  from 
the  molecular  weights  of  the  dissolved  substance  and  the  solvent, 
to  calculate  how  many  molecules  of  water  were  combined  with  one 
molecule  of  the  dissolved  substance.  The  results  of  such  a  calculation 
are  only  approximations.  In  the  first  place,  the  conductivity  method 
of  measuring  dissociation  is  not  accurate  for  concentrated  solutions, 
and  there  is  no  thoroughly  accurate  method  known  for  this  purpose. 
The  error  here  is,  however,  in  all  probability  not  very  large.  Another 
source  of  error,  which  is  probably  larger,  results  from  the  assumption 
that  Raoult's  law  holds  for  concentrated  solutions,  i.  e.,  that  for  con- 
centrated solutions  the  lowering  of  the  freezing-point  is  proportional 
to  the  concentration.  This  is  not  strictly  the  case,  and  we  do  not 
know  at  present  how  wide  the  deviation  from  Raoult's  law  is  in  con- 
centrated solutions. 

Carnegie  last.  Wash.  Pub.  No.  60. 


146  DISCUSSION   OF  EVIDENCE. 

Taking  all  of  these  factors  into  account,  it  still  seems  highly  probable 
that,  by  the  method  outlined  above,  we  can  arrive  at  a  reasonably 
close  approximation  to  the  amount  of  water  combined  with  a  molecule, 
or  the  resulting  ions,  of  a  dissolved  substance,  under  given  conditions 
of  concentration.  Whatever  objection  may  be  offered  to  this  method 
of  calculating  the  approximate  composition  of  the  hydrates  existing 
in  aqueous  solution,  it  should  be  stated  that  it  is  the  only  general 
method  thus  far  worked  out  for  throwing  any  light  whatever  on 
this  important  problem.  Jones  and  Getman  applied  this  method  of 
calculating  the  approximate  composition  of  hydrates  to  about  100 
compounds — salts,  acids,  and  organic  substances — and  to  about  1,500 
solutions  of  these  substances.  Their  results  have  been  recorded  in 
Publication  No.  60  of  the  Carnegie  Institution  of  Washington. 

Salts  of  lithium  form  more  complex  hydrates  than  those  of  sodium 
and  potassium.  This  would  be  expected,  since  lithium  salts  crystallize 
with  water,  while  the  salts  of  the  other  alkalies  in  general  crystallize 
without  water. 

Salts  of  potassium  and  ammonium  generally  crystallize  without 
water,  and  these  compounds,  as  would  be  expected,  combine  with  rela- 
tively little  water  in  aqueous  solution. 

Many  salts  of  sodium  crystallize  without  water,  and  these  hydrate 
very  slightly.  Other  sodium  salts,  such  as  the  bromide  and  iodide,  crys- 
tallize with  water  and  show  considerable  hydrating  power  in  solution. 

Salts  of  calcium  crystallize  with  water  and  all  have,  as  would  be 
expected,  large  hydrating  power.  The  halogen  salts  crystallize  with 
6,  the  nitrate  with  4  molecules  of  water.  The  nitrate  was  found  to 
have  less  hydrating  power  than  the  chloride  or  bromide. 

The  salts  of  strontium  resemble  those  of  calcium,  both  in  the  amounts 
of  water  with  which  they  crystallize  and  with  which  they  combine  in 
aqueous  solution.  Salts  of  barium  crystallize  with  less  water  and 
show  less  hydration  than  those  of  calcium  and  strontium. 

The  salts  of  magnesium  have  just  about  the  hydration  that  would 
be  expected  from  their  water  of  crystallization.  The  same  may  be 
said  of  the  salt  of  zinc  that  was  studied. 

Cadmium  is  of  special  interest.  Its  halogen  compounds  crystallize 
with  little  or  no  water,  and  although  cadmium  belongs  in  the  same 
group  with  metals  of  large  hydrating  power,  its  halogens  combine 
with  only  a  small  amount  of  water.  The  nitrate  of  cadmium  crystal- 
lizes with  4  molecules  of  water  and,  as  could  be  predicted,  shows  con- 
siderable hydrating  power. 

The  chloride  and  nitrate  of  magnesium  show  the  hydration  that 
would  be  expected  from  their  water  of  crystallization.  The  same  may 
be  said  of  the  salts  of  nickel,  cobalt,  and  copper. 

The  chlorides  and  nitrates  of  aluminium,  iron,  and  chromium  crys- 
tallize with  large  amounts  of  water  and  show  great  hydrating  power. 


DISCUSSION    OF   EVIDENCE.  147 

The  strong  mineral  acids  show  some  hydrating  power,  but  the  com- 
plexity of  the  hydrate  formed  by  these  substances  seems  to  pass 
through  a  maximum.  The  acids  thus  differ  from  the  salts. 

Some  13  non-electrolytes  were  studied  as  to  their  hydration,  and 
none  of  them  showed  any  appreciable  hydration.  The  same  applies 
to  the  organic  acids  that  were  studied  in  this  connection. 

The  following  general  relations  were  brought  out  by  the  work  of 
Jones  and  Getman.  The  total  amount  of  water  held  in  combination  by 
the  dissolved  substance  increases  as  the  concentration  of  the  solution 
increases.  From  what  is  known  of  mass  action,  this  would  be  expected. 

The  number  of  molecules  of  water  combined  with  1  molecule  of  the 
dissolved  substance  generally  increases  from  the  most  concentrated  to 
the  most  dilute  solution  studied.  In  some  cases,  however,  the  number 
of  molecules  of  combined  water  seems  to  pass  through  a  maximum. 
These  results,  we  believe,  give  us  the  approximate  amounts  of  com- 
bined water,  and  certainly  the  relative  hydrating  powers  of  the  different 
compounds  with  which  we  worked. 

One  other  relation  bearing  on  the  question  of  hydration  in  aqueous 
solution  was  brought  out  by  the  work  of  Jones  and  Getman. 

RELATION  BETWEEN  THE  MINIMA  IN  THE  FREEZING-POINT  CURVES  AND 
THE  MINIMA  IN  THE  BOILING-POINT  CURVES. 

It  has  already  been  pointed  out  that  if  we  plot  the  molecular  lower- 
ings  of  the  freezing-point  as  ordinates  against  the  concentrations  of 
the  solutions  as  abscissae,  the  curves  have  a  well-defined  minimum; 
from  this  minimum  they  rise  both  with  dilution  and  with  concentration. 
What  is  the  meaning  of  this  minimum?  The  curves  rise  with  dilution 
because  of  increasing  dissociation ;  they  rise  with  increasing  concentration 
because  the  total  amount  of  combined  water  increases  with  the  concen- 
tration of  the  solution;  and  consequently  the  lowering  of  the  freezing- 
point  of  the  ever-decreasing  amount  of  solvent  water  becomes  greater  and 
greater.  The  minima  in  freezing-point  curves  are,  then,  the  points  where 
the  two  opposite  effects  increase  in  dissociation  with  dilution  and  in- 
crease in  combined  water  with  concentration,  just  offset  one  another. 

The  rise  in  the  boiling-points  of  solvents  produced  by  dissolved  sub- 
stances was  also  studied,  and  at  different  concentrations.  Boiling-point 
curves  were  plotted  analogous  to  the  freezing-point  curves;  i.  e.,  molec- 
ular rise  in  the  boiling-point  as  ordinates  and  concentrations  as  abscissae. 
These  curves  also  had  minima,  and  we  interpreted  the  minima  here  in 
a  manner  analogous  to  our  interpretation  of  the  minima  in  the  freezing- 
point  curves — they  are  the  points  of  equilibrium  between  increasing  dis- 
sociation with  dilution,  and  increasing  hydration  with  concentration. 

A  comparison  of  the  freezing-point  with  the  boiling-point  curves 
brought  out  a  relation  of  interest,  and  we  believe  of  some  importance 
in  the  present  connection.  The  minima  in  the  boiling-point  curves 


148  DISCUSSION   OF   EVIDENCE. 

occur  at  greater  concentration  than  in  the  freezing-point  curves. 
What  is  the  meaning  of  this  relation?  To  see  this  we  must  first  call 
attention  to  one  property  of  hydrates  which  has  thus  far  not  been 
referred  to.  They  are  very  unstable,  and  readily  break  down  with 
rise  in  temperature.  This  is  easily  seen  if  we  consider  the  facts  in  the 
case  of  a  salt  like  calcium  chloride.  At  ordinary  temperatures  there 
may  be  as  much  as  30  molecules  of  water  combined  with  1  molecule 
of  the  salt;  while  at  the  boiling-point  of  the  saturated  solution  all  of 
the  water  can  be  removed  except  the  6  molecules  with  which  the  salt 
crystallizes.  The  higher  the  temperature  to  which  a  solution  is  heated 
the  less  the  hydration  in  such  a  solution.  Solutions,  of  course,  boil 
higher  than  they  freeze.  There  is  therefore  less  hydration  in  the 
boiling  than  in  the  freezing  solution.  Consequently,  to  produce  enough 
hydration  to  give  the  minimum  in  the  curve  would  require  a  greater 
concentration  of  the  solution  at  its  boiling-point  than  at  its  freezing-point . 
The  fact  is,  then,  not  only  in  accord  with  the  hydrate  conception, 
but  could  readily  have  been  predicted  from  it,  as  a  necessary  conse- 
quence of  the  theory. 

RELATION  BETWEEN  WATER  OF  CRYSTALLIZATION  AND  TEMPERATURE  OF 
CRYSTALLIZATION. 

Jones  and  Bassett1  worked  out  the  approximate  composition  of  the 
hydrates  formed  by  a  large  number  of  substances,  and  also  the  following 
relation.  The  hydrates,  as  we  have  seen,  are  very  unstable  systems. 
They  are  readily  broken  down  in  solution  with  rise  in  temperature. 
The  hydrates  which  exist  in  solution  at  ordinary  temperatures  are 
much  more  complex  than  those  which  the  salts  can  bring  with  them 
out  of  solution  as  water  of  crystallization.  The  hydrates  are  more 
stable  and  more  complex  the  lower  the  temperatures.  We  were,  how- 
ever, surprised,  on  examining  the  literature,  to  find  the  large  number 
of  examples  on  record  of  salts  crystallizing  with  varying  amounts  of 
water,  depending  on  the  temperature  at  which  the  crystals  were  formed. 

A  few  examples  will  be  given  to  bring  out  the  general  relation  that 
the  number  of  molecules  of  water  of  crystallization  is  larger  the  lower 
the  temperature  at  which  the  salt  crystallizes. 


CaCl2 

o  w2r»      -A-3  ^6  temperature  of  crys- 
A  S2X  \       tallization   is   lower   and 

4  H2<->               1_ 

MgCl2      6  H2O  1   Elevated  temperatures 
8  H2O  /       above  20°. 
10H20  20°. 

6  H2O  J       """"' 

12H2O.... 

—10°  to  —12°. 

MnCl2 

2H20  

....     20°. 

MnSO4   3H2O.... 

25°  to  31°,  as 

a 

4H2O  

....      15°. 

4H2O.... 

25°  to  31°. 

6H2O  

....-21°. 

5H2O.... 

15°  to  20°. 

11  H2O  

....-21°  to  -37°. 

7HjO.... 

0°  or  below  0°. 

12H2O  

.  .  .  .  -48°. 

FeCls 

anhydrous.  .  .  . 

80°  and  above. 

FeCls     3JH2O  

20°. 

2H2O  

.  .  .  .     60°  to  80°. 

6    H2O  

20°  to  -  16°. 

2|H20  

.  .  .  .     40°  to  60°. 

Carnegie  Inst.  Wash.  Pub.  No.  60.     Amer.  Chem.  Journ.,  33,  534  (1905) ;  34,  291  (1905).     Zeit. 
phys.  Chem.,  52,  231  (1905). 


DISCUSSION   OF   EVIDENCE.  149 

These  examples  suffice  to  show  the  general  nature  of  the  relation 
between  water  of  crystallization  and  the  temperature  at  which  the 
salt  crystallizes.  This  relation  could  have  been  foreseen  as  a  necessary 
consequence  of  the  theory  of  hydrates  in  aqueous  solution,  and  the 
instability,  with  temperature,  of  those  hydrates. 

HYDRATE  THEORY  IN  AQUEOUS  SOLUTIONS  BECOMES  THE  SOLVATE  THEORY  IN 
SOLUTIONS  IN  GENERAL. 

The  earliest  work  on  the  problem  of  the  nature  of  solution  was 
limited  to  water  as  the  solvent.  It  was  found  that  salts  in  general 
have  the  power  to  combine  with  more  or  less  of  the  water  in  which  they 
are  dissolved — have  a  greater  or  less  hydrating  power.  This  power 
is,  however,  possessed  to  a  very  different  degree  by  the  different  com- 
pounds. As  we  have  seen,  the  degree  of  hydration  of  a  salt  can  be 
approximately  determined  by  the  amount  of  water  with  which  it 
crystallizes  at  ordinary  temperatures. 

It  having  been  made  probable  that  hydration  exists  in  aqueous 
solution,  the  question  arose,  do  dissolved  substances  have  the  power 
to  combine  with  other  solvents  in  which  they  are  dissolved? 

To  test  this  Jones  and  Getman1  studied,  by  the  boiling-point  method, 
solutions  of  lithium  chloride  and  nitrate,  and  calcium  nitrate  in  ethyl 
alcohol.  They  used  also  a  number  of  other  salts.  It  was  found  that 
the  molecular  rise  in  the  boiling-point  was  not  only  greater  than  the 
theoretical  rise  at  nearly  all  of  the  concentrations  studied,  but  the 
molecular  rise  increases  rapidly  with  the  concentration  of  the  solution. 
The  molecular  rise  of  the  boiling-point  of  ethyl  alcohol,  produced  by 
lithium  chloride,  increases  from  1.28  at  0.07  normal  to  2.43  at  2.07 
normal.  In  calculating  the  theoretical  molecular  rise  the  dissociation 
is,  of  course,  taken  into  account.  The  dissociation  decreases  with  the 
concentration,  which  would  tend  to  decrease  the  molecular  rise  in  the 
boiling-point.  Notwithstanding  this  influence,  we  have  seen  that  the 
molecular  rise  in  the  boiling-point  of  solutions  of  certain  salts  in  ethyl 
alcohol  increases  as  the  concentration  of  the  solution  increases. 

The  differences  between  the  theoretical  and  the  experimental  results 
are  in  some  cases  quite  large.  Jones  and  Getman2  interpreted  these 
results  in  ethyl  alcohol  in  a  manner  analogous  to  that  which  they  had 
adopted  in  the  case  of  aqueous  solutions.  The  abnormally  large  rise 
in  the  boiling-point  of  ethyl  alcohol,  produced  by  certain  salts,  and 
the  increase  in  the  molecular  rise  of  the  boiling-point  with  increase 
in  the  concentration  of  the  solution,  are  due  to  combination  between 
the  dissolved  substance  and  part  of  the  solvent — to  the  formation  of 
ethyl  alcoholates  in  solution.  The  part  of  the  alcohol  that  is  combined 
with  the  dissolved  substance  is  thus  removed  from  the  field  of  action 
as  far  as  solvent  is  concerned.  There  being  less  alcohol  present  acting 

1Amer.  Chem.  Journ.,  32,  338  (1904).  *Ibid.,  339  (1904). 


150  DISCUSSION   OF   EVIDENCE. 

as  solvent,  the  rise  in  its  boiling-point  produced  by  a  given  amount 
of  dissolved  substance  would  be  larger  than  if  all  the  alcohol  were 
playing  the  part  of  solvent. 

Further,  if  the  dissolved  substance  combines  with  a  part  of  the 
alcohol,  the  more  concentrated  the  solution  the  greater  the  total 
amount  of  alcohol  held  in  combination.  This  would  explain  the 
increase  in  the  molecular  rise  in  the  boiling-point  with  increase  in  tha 
concentration  of  the  solution.  This  suggestion  of  combination  between 
a  part  of  the  solvent  and  the  dissolved  substance  explains  the  facts 
in  alcoholic  solutions  just  as  well  as  the  hydrate  theory  explains  the 
facts  in  aqueous  solutions. 

The  work  of  Jones  and  Getman  in  solutions  in  ethyl  alcohol  as  the 
solvent  was  extended  by  Jones  and  McMaster  to  methyl  alcohol. 
They  also  extended  the  work  in  ethyl  alcohol  as  the  solvent.  They 
repeated  a  part  of  the  work  of  Jones  and  Getman  in  ethyl  alcohol  and 
obtained  results  of  the  same  general  character  as  had  been  found  by 
the  earlier  workers. 

They  used  the  boiling-point  method  with  methyl  alcohol  as  the 
solvent,  and  the  chloride,  bromide,  and  nitrate  of  lithium  as  the  dis- 
solved salts.  The  molecular  rise  in  the  boiling-point,  even  in  the  most 
dilute  solutions,  was  greater  than  could  be  accounted  for  by  the  disso- 
ciation. This  is,  of  course,  entirely  incapable  of  accounting  for  the 
increase  in  the  molecular  rise  with  increase  in  the  concentration  of  the 
solution,  which  manifests  itself  in  the  case  of  every  salt  studied  in  this 
solvent,  dissociation  decreasing  with  increase  in  concentration,  which 
would  tend  to  diminish  the  molecular  rise  in  the  boiling-point. 

The  magnitude  of  the  molecular  rise  in  the  most  concentrated  solu- 
tions is  very  large  indeed.  It  is  almost  twice  the  boiling-point  constant, 
or  normal  molecular  rise  for  this  solvent;  and  the  dissociation  of  such 
solutions  is  certainly  not  more  than  25  to  30  per  cent,  and  probably 
less  than  this  value. 

We  interpreted  these  results  as  we  did  those  in  ethyl  alcohol  as  the 
solvent — there  is  combination  between  a  part  of  the  alcohol  present 
and  the  dissolved  substance,  forming  methyl  alcoholates.  As  the  con- 
centration increases,  more  and  more  alcohol  is  held  in  combination 
by  the  dissolved  substance;  consequently,  there  is  an  increase  in  the 
molecular  rise  of  the  boiling-point. 

It  thus  seems  that  evidence  was  furnished  of  combination  between 
methyl  alcohol  and  the  dissolved  substance,  on  the  one  hand,  and 
ethyl  alcohol  and  the  dissolved  substance  on  the  other.  As  we  shall 
see  later,  evidence  has  been  obtained  of  combination  between  acetone 
and  substances  dissolved  in  it;  and  other  solvents  have  been  and  are 
being  brought  within  the  scope  of  this  work. 

In  every  case  thus  far  investigated  there  seems  to  be  good  evidence 
in  favor  of  the  view  that  there  is  combination  between  the  dissolved 


DISCUSSION   OF   EVIDENCE.  151 

substance  and  a  part  of  the  solvent  present.  In  a  word,  combination 
of  solvent  with  dissolved  substance — solvation — seems  to  be  a  more  or 
less  general  phenomenon.  The  original  hydrate  theory  thus  becomes 
the  solvate  theory  of  solution. 

TEMPERATURE  COEFFICIENTS  OF  CONDUCTIVITY  AND  HYDRATION. 

A  fairly  elaborate  investigation  on  the  conductivities,  dissociation, 
and  temperature  coefficients  of  conductivity  and  dissociation  of  aque- 
ous solutions  was  begun  in  my  laboratory  about  15  years  ago  and  is 
still  in  progress.  The  work,  as  a  whole,  has  been  recently  published 
by  the  Carnegie  Institution  of  Washington.1  The  monograph  in 
question  contains  the  investigations  of  Clover,2  Hosford,3  Howard,4 
Jacobson,5  Kreider,6  Shaeffer,7  Smith,8  Springer,9  West,10  Wight,11 
Wightman,12  and  Winston.13  The  results  published  in  this  monograph 
are  for  about  110  salts,  which  were  studied  from  zero  to  65°,  and  from 
the  most  concentrated  solution  that  could  be  used  to  the  dilution 
in  most  cases  of  complete  dissociation.  The  temperature  coefficients 
of  conductivity  were  calculated  both  in  conductivity  units  and  in 
percentage. 

Similar  data  were  obtained  for  about  90  of  the  more  common  organic 
acids,  and  the  constants  for  the  weaker  acids  were  calculated  from  the 
Ostwald  dilution  law.  The  dissociations  of  the  salts  and  acids  at  the 
different  temperatures  were,  whenever  possible,  also  calculated. 

The  temperature  coefficients  of  conductivity  were  calculated  both 
in  percentage  and  in  conductivity  units.  A  study  of  the  temperature 
coefficients  of  conductivity,  expressed  in  conductivity  units,  brought 
out  a  relation  which  had  a  very  direct  bearing  on  the  question  of 
hydration  in  aqueous  solution.  This  is  so  important  that  it  will  be 
discussed  here  in  some  detail. 

The  conductivity  of  a  solution  is  conditioned  by  the  number  of  ions 
present  and  the  velocities  with  which  they  move.  Rise  in  temperature 
not  only  does  not  increase  the  number  of  ions  present,  but,  as  is  well- 
known,  diminishes  dissociation.  The  effect  of  rise  in  temperature 
increasing  the  conductivity  of  solutions  is,  then,  due  to  an  increase  in 
the  velocities  with  which  the  ions  move.  If  the  ion  is  driven  by  a 
constant  force,  its  velocity  would  be  determined  chiefly  by  the  viscosity 
of  the  solvent  and  by  the  mass  and  size  of  the  ion.  With  rise  in 
temperature  the  driving  force  would  be  increased.  Rise  in  tempera- 
ture would  also  decrease  the  viscosity  of  the  solvent.  The  effect  of 

Carnegie  Inst.  Wash.  Pub.  No.  170.  *Ibid.,  50,  1  (1913). 

2Amer.  Chem.  Journ.,  43,  187  (1910).         9Ibid.,  48,  411  (1912). 

3Ibid.,  46,  240  (1911).  "Ibid.,  44,  508  (1910). 

*Ibid.,  48  500  (1912).  "Ibid.,  42,  520  (1909);  44,  159  (1910). 

&Ibid.,  40  355  (1908).  albid.,  46,  56  (1911);  48,  320  (1912). 

•Ibid.,  45,  282  (1911).  "Ibid.,  46,  368  (1911). 

Ubid.,  49,  207  (1913). 


152 


DISCUSSION   OF   EVIDENCE. 


rise  in  temperature  on  both  of  these  factors  would  be  to  increase  the 
velocities  of  the  ions  and,  consequently,  the  conductivity. 

Another  factor  must,  however,  be  taken  into  account.  That  many 
ions  in  aqueous  solutions  are  strongly  hydrated  seems  now  quite 
generally  accepted.  We  have  seen  that  these  hydrates  are  relatively 
unstable  and  break  down  with  rise  in  temperature.  The  simpler  the 
hydrate  formed  by  an  ion,  the  smaller  the  mass  of  the  ion;  the  smaller 
the  mass  of  the  ion,  other  things  being  equal,  the  less  resistance  it 
will  offer  when  moving  through  the  solvent.  Therefore,  rise  in  tem- 
perature should  increase  the  velocity  of  the  ion. 

TABLE  48. — Temperature  coefficients  of  conductivity. 


Temperature  coefficients  ia 

conductivity  units. 

hydrating  power. 

25°  to  35° 

50°  to  65° 

0  =  8 

»  =  1024 

v  =  8 

o=1024 

Sodium  chloride  

2.00 

2.46 

2.27 

2.82 

Sodium  bromide  

1.89 

2.18 

2.79 

Sodium  iodide  

2.12 

2.54 

2.33 

3.14 

Sodium  nitrate  

2.04 

2.45 

2.02 

2.67 

Sodium  chlorate  

1.77 

2.22 

2.15 

2.90 

Potassium  chloride  

2.39 

2.84 

2.45 

3.11 

Potassium  bromide  

2.43 

2.91 

2.45 

3.11 

Potassium  iodide  

2.38 

2.91 

2.65 

3.37 

Potassium  nitrate  

2.08 

2.16 

2.31 

2.83 

Potassium  chlorate  

2.02 

2.52 

2.23 

2.94 

Potassium  permanganate.  .  . 

2.04 

2.31 

2.29 

2.23 

Potassium  sulphocyanate.  .  . 

2.20 

2.56 

2.34 

Ammonium  chloride  

2.42 

2.94 

2.51 

3.69 

Ammonium  bromide  

2.32 

2.86 

2.58 

3.11 

Ammonium  nitrate  

2.17 

2.50 

2.33 

3.04 

If  decreasing  complexity  of  the  hydrate  formed  by  the  ion  with 
rise  in  temperature  plays  any  prominent  part  in  determining  the  large 
temperature  coefficient  of  conductivity,  since  the  complexity  of  such 
hydrates  would  decrease  more  with  rise  in  temperature,  we  should 
expect  to  find  that  the  ions  which  have  the  greatest  hydrating  power 
would  have  the  largest  temperature  coefficients  of  conductivity.  This 
is  a  concrete  and,  it  would  seem,  necessary  consequence  of  the  hydrate 
theory  in  aqueous  solutions.  Further,  it  is  one  which  can  be  tested 
directly  by  experiment.  Is  it  true? 

We  have  seen  that  the  hydrating  power  of  a  salt,  or  the  ions  into 
which  it  dissociates,  is  approximately  proportional  to  the  number  of 
molecules  of  water  with  which  it  crystallizes.  This  is  the  same  as  to 
say  that  the  salt  which  has  the  greatest  power  to  bring  water  with  it 
out  of  solution  is  the  one,  other  things  being  equal,  which  would  hold 
the  largest  number  of  molecules  of  water  in  combination  with  it  in 
solution.  The  question  is,  therefore,  is  there  any  relation  between 


DISCUSSION    OF   EVIDENCE. 


153 


the  number  of  molecules  of  water  with  which  a  salt  crystallizes  and 
its  temperature  coefficients  of  conductivity? 

This  relation  has  already  been  discussed  in  Publication  No.  170  of 
the  Carnegie  Institution  of  Washington,  in  which  the  results  of  our 
work  on  conductivity  and  dissociation  has  been  published.  Tables  48 
and  49,  showing  temperature  coefficients  in  conductivity  units  between 
the  temperatures  25°  and  35°,  on  the  one  hand,  and  between  50°  and 
65°  on  the  other,  at  the  dilutions  |  and  y^  normal,  are  taken  from  the 
monograph  referred  to  above. 

TABLE  49. — Temperature  coefficients  of  conductivity. 


Temperature  coefficients  in 

conductivity  units. 

Substances  with  large 
hydrating  power. 

25°  to  35° 

50°  to  65° 

»«8 

r  =  1024 

«  =  8 

0  =  1024 

Calcium  chloride  

3.49 

4.85 

Calcium  bromide 

3.73 

5.00 

4.03 

6  03 

Calcium  nitrate 

3.09 

4.79 

3.33 

Strontium  chloride  

3.37 

5^13 

3.92 

6.02 

Strontium  bromide  

3.66 

5.27 

4.08 

Strontium  nitrate  

2.76 

4.86 

3.58 

Barium  nitrate  

3.09 

4.74 

3.34 

Magnesium  chloride  

3.40 

4.72 

3.61 

Magnesium  bromide  

3.55 

4.44 

4.08 

Magnesium  nitrate  

3.10 

4.78 

3.57 

Zinc  nitrate  

3.13 

4.47 

3.43 

5.41 

Manganous  chloride  

3.14 

4.86 

3.43 

6.37 

Nickel  chloride  

3.41 

5.04 

3.01 

Nickel  nitrate  

3.21 

4.58 

Cobalt  chloride  

3.39 

4.95 

3.54 

Cobalt  bromide  

3.32 

4.96 

3.75 

Cobalt  nitrate  

3.20 

4.67 

3.05 

Cupric  nitrate  

3.18 

4.88 

Aluminium  chloride  

4.57 

8.64 

5.16 

12.49 

Aluminium  nitrate  

4.19 

7.86 

4.87 

11.65 

We  have  seen  that  the  hydrates  formed  by  a  large  number  of  salts, 
including  those  given  in  tables  48  and  49,  have  already  been  worked 
out,1  and  that  water  of  crystallization  is  a  rough  measure  of  water  of 
hydration.  The  salts  in  table  48  crystallize  with  little  or  no  water, 
and  in  aqueous  solution  are  very  little  hydrated;  those  in  table  49,  in 
general,  crystallize  with  large  amounts  of  water  and  are  strongly 
hydrated  compounds. 

Let  us  compare  the  temperature  coefficients  of  conductivity  in  con- 
ductivity units  (which  are  the  actual  increases  in  molecular  conduc- 
tivity per  degree  rise  in  temperature)  of  the  substances  in  table  48 
with  those  in  table  49.  It  will  be  seen  that  the  coefficients  for  the  sub- 
stances in  table  48  are,  at  all  dilutions  and  temperatures,  much  smaller 

'Carnegie  Inst.  Wash.  Pub.  No.  60. 


154  DISCUSSION   OF   EVIDENCE. 

than  those  in  table  49.  In  making  this  comparison  we  must,  of  course, 
take  into  account  the  fact  that  the  substances  in  table  48  are  binary 
electrolytes,  each  molecule  breaking  down  into  2  ions,  while  the  sub- 
stances recorded  in  table  49  are  all  ternary  electrolytes,  each  mole- 
cule breaking  down  into  3  ions,  except  the  two  salts  of  aluminium 
which  are  quaternary  electrolytes,  each  molecule  yielding  4  ions.  Even 
taking  all  of  these  facts  into  account,  the  temperature  coefficients 
of  conductivity  for  the  slightly  hydrated  salts  are  much  smaller  than 
those  for  the  strongly  hydrated  compounds.  This  is  exactly  what 
would  be  expected.  The  complexity  of  the  hydrates  of  slightly 
hydrated  salts  could  change  only  a  little  with  rise  in  temperature. 
Consequently,  the  mass  of  the  hydrated  ion  would  change  only  slightly 
with  rise  in  temperature,  and  this  effect  of  temperature  on  conductivity 
would  be  very  small. 

Another  relation  manifests  itself  when  we  compare  the  results  in 
table  48  with  one  another,  and  those  in  table  49  with  one  another.  If 
the  temperature  coefficient  of  conductivity  is  a  function  of  the  decreas- 
ing complexity  of  the  hydrate  formed  by  the  ion,  as  the  temperature 
is  raised  we  should  expect  that  those  substances  which  have  equal 
hydrating  power  would  have  approximately  the  same  temperature 
coefficients  of  conductivity. 

The  substances  in  table  48  have  only  slight  hydrating  power,  shown 
by  the  fact  that  they  crystallize  with  little  or  no  water.  The  fact  is, 
their  temperature  coefficients  of  conductivity  are  all  of  the  same  order 
of  magnitude. 

The  salts  in  table  49  have  different  hydrating  power,  but  all  have 
great  power  to  combine  with  water  in  aqueous  solution.  A  large 
number  of  these  compounds  have  approximately  the  same  hydrating 
power,  as  would  be  expected  from  the  fact  that  many  of  them  crystal- 
lize with  6  molecules  of  water.  Barium  chloride  crystallizes  with  only 
2  molecules  of  water,  yet  forms  hydrates  of  complexity  comparable 
with  the  other  salts1  in  this  table.  Its  temperature  coefficients  are  of 
the  same  order  of  magnitude  as  those  of  the  other  substances  in  the 
table.  Manganous  chloride  with  4  molecules  of  water  of  crystalliza- 
tion, and  copper  chloride  with  2,  form  hydrates  of  the  same  degree  of 
complexity  as  the  other  salts  in  this  table.  Their  temperature  coef- 
ficients are  in  keeping  with  this  fact.  The  chloride  of  aluminium 
crystallizes  with  6  molecules  of  water  and  the  nitrate  with  8.  These 
salts,  as  has  already  been  pointed  out,  break  down  yielding  4  ions 
each.  Their  temperature  coefficients  are  larger  than  those  of  the 
ternary  electrolytes. 

The  more  dilute  the  solution,  the  more  complex  the  hydrate  formed 
by  the  molecule  or  the  ion.  This  is  but  the  expression  of  the  action 

Carnegie  Inst.  Wash.  Pub.  No.  60,  pp.  75,  76. 


DISCUSSION   OF   EVIDENCE.  155 

of  mass:  the  more  water  there  is  present  the  more  will  be  combined 
with  the  dissolved  substance.  The  more  complex  the  hydrate  the 
greater  the  change  in  the  complexity  of  the  hydrate  with  rise  in  tem- 
perature. Since  the  magnitude  of  the  temperature  coefficients  of 
conductivity  seems  to  be  a  function  of  the  change  in  the  complexity 
of  the  hydrate  with  rise  in  temperature,  it  follows,  from  the  hydrate 
theory,  that  the  temperature  coefficients  of  conductivity  for  any  given 
substance  should  be  greater  at  the  higher  dilution  than  at  the  lower. 

A  comparison  of  the  results  at  the  two  dilutions  for  any  given  sub- 
stance in  table  48  or  table  49  will  show  that  the  above  consequence  of 
the  hydrate  theory  is  confirmed  by  the  facts.  The  temperature  coef- 
ficients are  larger  at  the  higher  dilution  for  every  substance  recorded 
in  both  tables. 

One  other  relation  should  be  pointed  out  before  leaving  the  discussion 
of  the  temperature  coefficients  of  conductivity.  We  have  seen  that 
the  hydrates  are  unstable,  and  that  with  rise  in  temperature  they 
break  down.  The  higher  the  temperature  to  which  they  are  heated 
the  more  unstable  they  become.  We  should,  therefore,  expect  the 
hydrates  to  break  down  more  rapidly  as  the  temperature  goes  higher. 
If  this  were  the  case,  the  higher  the  temperature  of  the  solution  the 
larger  the  temperature  coefficients  of  conductivity.  If  we  compare 
the  results  for  any  given  substance  in  table  48  or  49  we  will  find  that 
such  is  the  case.  The  temperature  coefficients  for  any  given  dilution 
are  higher  between  50°  and  65°  than  between  25°  and  35°. 

The  above  four  conclusions  from  the  solvate  theory  of  solution,  as 
far  as  aqueous  solutions  are  concerned,  are  confirmed  at  every  point 
by  the  results  of  measuring  the  temperature  coefficients  of  conductivity. 
Without  this  theory  it  does  not  appear  to  be  simple  to  explain  the 
above  relations.  The  agreement  between  the  four  deductions  from 
the  theory  and  the  experimental  results  is  so  satisfactory  that  it  is 
regarded  as  strong  evidence  in  favor  of  the  general  correctness  of 
the  theory. 

RELATION  BETWEEN  THE  HYDRATION  OF  THE  IONS  AND  THEIR  IONIC  VOLUMES. 

Jones  and  Pearce1  worked  out  the  approximate  composition  of  the 
hydrates  formed  by  a  large  number  of  salts,  using  the  freezing-point 
and  conductivity  methods  already  referred  to.  They  found  the 
following  relation  between  the  volumes  of  the  ions  and  their  power 
to  form  hydrates.  The  atomic  volume  curve  is  obtained  by  plotting 
the  atomic  weights  of  the  elements  as  abscissae  against  the  atomic 
volumes  as  ordinates.  This  curve,  as  is  well-known,  contains  well- 
defined  maxima  and  minima.  At  the  maxima  are  the  alkali  elements, 
the  three  with  the  largest  atomic  volumes  being  potassium,  rubidium, 

Carnegie  Inst.  Wash.  Pub.  No.  180,  p.  57;  Amer.  Chem.  Journ.,  38,  736  (1907). 


156  DISCUSSION   OF   EVIDENCE. 

and  caesium.  The  salts  of  these  elements  generally  crystallize  without 
water,  and  therefore  have  very  little  hydrating  power  in  aqueous 
solution.  The  approximate  hydration1  of  salts  of  potassium  has  been 
determined  by  the  method  usually  employed  and  has  been  found  to 
be  small. 

Some  of  the  salts  of  lithium  and  potassium  crystallize  with  2  and 
3  molecules  of  water,  and  these  have  been  shown  to  have  some  hydrat- 
ing power.1  The  atomic  volumes  of  lithium  and  sodium  are  much 
smaller  than  those  of  potassium,  rubidium,  and  cesium. 

Turning  from  the  maxima  of  the  curve  to  the  minima,  at  the  mini- 
mum of  the  third  section  of  the  curve  are  iron,  cobalt,  nickel,  and 
copper.  Salts  of  these  metals  crystallize  with  large  amounts  of  water, 
and  in  aqueous  solution  they  form  complex  hydrates. 

Aluminium  falls  at  the  second  minimum  of  the  atomic  volume  curve, 
having  a  somewhat  greater  atomic  volume  than  iron.  The  salts  of 
aluminium  crystallize  with  large  amounts  of  water,  some  of  them  with 
6  and  8  molecules.  In  aqueous  solution  they  form  complex  hydrates.1 
Barium  has  the  largest  atomic  volume  of  members  of  its  group;  its 
salts  crystallize  without  water  or  some  with  2  molecules  of  water. 
Many  of  the  salts  of  calcium,  strontium,  and  magnesium  crystallize 
with  6  molecules  of  water.  Magnesium  has  the  smallest  volume  of 
any  element  of  this  group;  it  has  been  found  to  have  the  greatest 
hydrating  power  of  any  member  of  the  group.  Strontium  has  a 
slightly  larger  atomic  volume  than  calcium  and  has  a  somewhat  smaller 
power  to  form  hydrates.  Taking  all  of  the  facts  into  account,  it 
would  seem  that,  other  things  being  equal,  the  smaller  the  cation  the 
greater  its  hydrating  power.  This  raises  the  question,  which  ion  is  it 
that  forms  the  hydrate?  Do  both  ions  form  hydrates.  If  so,  which 
has  the  greater  hydrating  power? 

The  different  salts  of  certain  metals  have  approximately  the  same 
hydrating  power.  The  common  constituent  of  these  salts  is  of  course 
the  cation,  the  anion  varying  from  salt  to  salt.  This  would  indicate 
that  it  is  primarily  the  cation  which  conditions  the  hydrating  power 
of  a  salt.  Since  the  different  salts  of  the  same  metal  do  not  all  have 
the  same  hydrating  power,  it  seems  reasonable  to  assume  that  the 
anion  has  some  power  to  form  hydrates  in  the  presence  of  water.  The 
cation  is,  then,  the  chief  hydrating  agent,  and  its  hydrating  power 
seems  to  be  a  function  of  its  size  or  atomic  volume — the  smaller  the 
ion  the  greater  its  power  to  hold  water  in  combination  with  it  in 
aqueous  solution. 

This  raises  the  question,  why  is  this  the  case?  It  has  occurred  to 
me  that  the  electrical  density  of  the  charge  on  the  ion  may  have  some- 
thing to  do  with  this  relation.  Other  things  being  equal,  the  smaller 

Carnegie  Inst.  Wash.  Pub.  No.  60. 


DISCUSSION   OF   EVIDENCE.  157 

ion  has  the  greater  density  of  charge  upon  its  surface;  this  might 
enable  it  to  hold  more  molecules  of  water  in  combination  with  itself. 
There  seem,  however,  to  be  certain  physical  objections  to  this  explana- 
tion of  the  relation  in  question.  Whatever  the  explanation,  the  fact 
remains. 

HYDRATION  OF  THE  IONS  AND  THE  VELOCITIES  WITH  WHICH  THEY  MOVE. 

Certain  apparent  discrepancies  presented  themselves  in  the  velocities 
of  the  different  ions,  which,  for  a  tune,  could  not  be  explained.  It  was 
known  that  the  lithium  ion,  under  the  same  driving  force,  moves  more 
slowly  than  potassium;  and  yet  it  has  smaller  volume  and  smaller 
mass.  It  was  not  until  it  was  shown1  that  the  lithium  ion  is  more 
strongly  hydrated  than  sodium  or  potassium  that  this  fact  could  be 
explained,  and  other  apparent  discrepancies  presented  themselves. 
A  relation  between  the  migration  velocities  of  the  ions  and  their  hydrat- 
ing  power  was  worked  out  by  Jones  and  Pearce.2  Their  discussion  is 
repeated  here  to  bring  out  the  point  in  question. 

The  velocities  of  the  ions  in  moving  through  any  given  medium  is 
known  to  vary  inversely  as  their  mass,  the  driving  force  being  constant. 
Their  velocities  would  also  vary  inversely  as  their  volumes.  Mass 
being  constant,  we  should  expect  the  ions  with  the  smallest  atomic 
volumes  to  move  the  swiftest  under  a  constant  driving  force,  while 
the  facts  are  often  the  opposite.  Leaving  out  of  account  the  hydrogen 
and  hydroxyl  ions,  potassium,  rubidium,  and  caesium  have  very  great 
velocities  and  the  largest  volumes;  while  the  ions  of  the  iron  and 
copper  group  have  the  smallest  volumes  and  very  small  velocities.  The 
meaning  of  this  apparent  discrepancy  can  be  seen  at  once  by  com- 
paring the  atomic  volume  curve  and  the  migration  velocity  curve. 

The  ions  with  the  smallest  volumes  have  the  greatest  hydrating 
power.  The  ions  with  the  smallest  volumes  frequently  have  the 
smallest  velocities.  Therefore,  the  ions  with  smallest  velocities  fre- 
quently have  the  greatest  hydrating  power.  To  discuss  the  relations 
somewhat  in  detail,  the  atomic  volumes  of  potassium,  rubidium,  and 
caesium  increase  rapidly  with  increasing  atomic  weight,  and  their  salts 
generally  crystallize  without  water.  The  atomic  volumes  of  sodium 
and  lithium  are  less  than  half  that  of  potassium,  and  yet  their  velocities 
are  only  about  two-thirds  that  of  potassium.  It  will  be  recalled  that 
salts  of  sodium  and  lithium  may  crystallize  with  2  or  3  molecules  of 
water.  We  may  therefore  assume  that  the  increase  in  the  volume 
and  mass  of  the  lithium  and  sodium  ions,  due  to  the  formation  of  a 
hydrate,  decreases  the  velocity  of  these  ions  below  that  of  potassium. 
The  small  velocity  of  the  lithium  ion  was,  as  we  have  seen,  for  a  long 
time  unexplained.  It  has  a  volume  only  about  half  that  of  sodium, 
and  the  largest  ascertained  amount  of  water  with  which  the  salts  of 

Carnegie  Inst.  Wash.  Pub.  No.  60.  'Ibid.,  180,  pp.  84-86. 


158  DISCUSSION   OF   EVIDENCE. 

lithium  crystallize  is  3.  The  maximum  amount  for  many  of  the  salts 
of  sodium  is  2.  The  lithium  ion  is,  in  general,  more  hydrated  than  the 
sodium  ion,  and  its  velocity  is  therefore  decreased  more  by  hydration. 
Notwithstanding  its  smaller  volume  and  lighter  mass,  on  account  of  its 
greater  hydration  lithium  moves  with  about  the  same  velocity  as 
sodium. 

The  calcium  ion  is  slightly  larger  than  sodium,  but  has  considerably 
smaller  velocity.  This  is  undoubtedly  due  primarily  to  its  much 
greater  hydrating  power.  Within  this  group  the  atomic  volumes 
increase  with  increasing  atomic  weight.  The  velocities  of  calcium 
and  strontium,  with  many  salts  crystallizing  with  6  molecules  of  water, 
are  approximately  equal  to  that  of  barium.  Many  of  the  salts  of 
barium  crystallize  with  2  molecules  of  water  or  water-free.  The  larger 
mass  of  the  barium  ion  itself  diminishes  the  velocity.  Magnesium, 
with  about  half  the  volume  of  calcium,  has  nearly  the  same  velocity, 
due  to  its  greater  hydrating  power.  The  cobalt,  nickel,  and  copper  ions 
have  nearly  the  same  volumes  and  approximately  the  same  hydrating 
power.  They  have  approximately  the  same  velocities. 

The  atomic  volumes  of  the  chloride,  bromide,  and  iodide  ions  are 
approximately  the  same.  If  they  hydrate  at  all  we  should  expect  the 
same  order  of  hydration  for  all  three,  as  has  been  made  probable.  We 
should  expect  them  to  have  velocities  of  the  same  order  of  magnitude, 
and  such  is  the  fact. 

The  silver  ion  is  the  only  well-established  exception.  It  has  a  small 
volume  and  many  of  its  salts  crystallize  without  water.  Although 
it  has  small  volume,  it  apparently  has  but  little  hydrating  power. 
Notwithstanding  its  considerable  mass,  with  its  small  volume  and 
small  hydrating  power  we  should  expect  it  to  have  a  fairly  high  velocity. 
The  fact  is,  the  velocity  of  the  silver  ion  is  slightly  less  than  that  of 
chlorine,  bromine,  and  iodine. 

The  general  truth  of  the  relation  that  the  ions  with  the  smallest 
velocities  have  the  greatest  hydrating  power  is,  then,  established  by 
the  facts,  the  great  hydrating  power  being  one  of  the  factors  condition- 
ing the  small  velocity. 

DISSOCIATION  AS  MEASURED  BY  THE  FREEZING-POINT  METHOD  AND  BY 
THE  CONDUCTIVITY  METHOD. 

When  the  theory  of  electrolytic  dissociation  was  proposed,  it  became 
a  problem  to  measure  accurately  the  magnitude  of  dissociation. 
Arrhenius  pointed  out,  in  his  original  epoch-making  paper,  that 
dissociation  could  be  measured  either  by  the  freezing-point  or  by  the 
conductivity  method.  The  conductivity  of  a  few  electrolytes  was  first 
worked  out  accurately  by  Friederich  Kohlrausch,  by  a  method  which 
he  devised  for  the  purpose.  This  work  was  done  before  the  theory 
of  electrolytic  dissociation  was  proposed.  Kohlrausch's  data  were 


DISCUSSION   OF   EVIDENCE.  159 

used  to  calculate  the  magnitude  of  the  dissociation  of  the  electrolytes 
with  which  he  worked,  at  the  various  dilutions  of  the  solutions. 

Another  method  of  measuring  the  dissociation  of  electrolytes,  based 
upon  the  change  in  the  solubility  of  a  salt  on  the  addition  of  a  second 
salt  with  a  common  ion,  was  developed  theoretically  by  Nernst1  while 
working  in  Ostwald's  laboratory.  When  applied  experimentally  it 
gave  dissociations  which  were  different  from  those  obtained  by  the 
conductivity  method.  The  freezing-point  method  had  not  been  used 
at  that  time  to  measure  dissociation.  There  was  then  a  period  when 
there  were  but  two  methods  for  measuring  dissociation,  and  these  gave 
widely  different  results.  During  this  period  the  point  was  made 
against  the  dissociation  theory,  that  if  dissociation  took  place  in  the 
presence  of  water  there  was  no  means  of  determining  its  magnitude. 

At  this  time  Ostwald  so  improved  the  freezing-point  method  that  it 
could  be  used  to  measure  dissociation.  He  started  me  to  work  on  the 
application  of  this  method,  and  we2  measured  the  dissociation  of  a  fairly 
large  number  of  salts.  The  results  differed  radically  from  those 
obtained  by  the  solubility  method,  but  agreed  fairly  well  with  those 
calculated  from  the  conductivity  method  of  Kohlrausch.  It  was 
afterwards  shown  that  an  assumption  had  been  made  in  applying  the 
solubility  method,  which,  when  corrected,  enabled  that  method  to  give 
essentially  the  same  results  as  those  obtained  by  the  other  two. 

A  comparison  of  the  data  from  the  freezing-point  method  with  those 
from  the  conductivity  method  showed  that  Association  as  measured 
by  the  former  was  slightly  higher  than  by  the  latter.  The  meaning 
of  this  discrepancy  was  at  that  time  not  understood. 

After  it  had  been  established,  with  reasonable  certainty,  that  hydra- 
tion  takes  place  in  aqueous  solution,  a  possible  explanation  of  this 
apparent  discrepancy  presented  itself.  But  before  offering  this  expla- 
nation it  seemed  desirable  to  do  more  experimental  work,  having  this 
point  especially  in  mind.  Pearce3  carried  out  in  my  laboratory  a  very 
careful  piece  of  work,  in  which  dissociation  was  measured  by  the 
freezing-point  method  and  also  by  the  conductivity  method,  and  the 
two  sets  of  results  were  compared.  We  worked  with  the  chlorides  of 
calcium,  strontium,  magnesium,  barium,  cobalt,  copper,  and  alumin- 
ium; with  the  nitrates  of  calcium,  magnesium,  barium,  cobalt,  nickel, 
and  copper;  with  sodium  bromide,  and  with  hydrochloric,  nitric,  and 
sulphuric  acids. 

That  hydration  can  explain  the  fact  that  dissociation  as  measured 
by  freezing-points  is  higher  than  as  measured  by  conductivity  can  be 
seen  from  the  following.  The  combined  water  is  removed  from  the 
field  of  action  as  solvent ;  only  the  uncombined  water  is  acting  as 
solvent.  Freezing-point  lowering  is  proportional  to  the  ratio  between 

^eit.  phys.  Chem.,  4,  1372  (1889).  'Carnegie  Inst.  Wash.  Pub.  No.  180,  p.  57. 

2Ibid.,ll,  HO,  529;  12,  633  (1893). 


160 


DISCUSSION   OF   EVIDENCE. 


the  number  of  molecules  of  the  dissolved  substance  and  of  the  solvent. 
If  one-fourth  of  the  water  present  is  combined  with  the  dissolved 
substance,  the  freezing-point  lowering  would  be  one-fourth  greater 
than  if  all  the  water  were  present  as  free  water  and  therefore  acting 
as  solvent  water.  Freezing-point  lowering  would  thus  be  affected  pro- 
portionally by  hydration.  Dissociation  of  concentrated  solutions 
calculated  from  the  freezing-point  lowering  would  therefore  be  much 
too  high. 

The  conductivity  of  a  solution  depends  upon  the  number  of  ions 
present  and  their  velocities.  The  number  of  ions  would  probably  not 
be  affected  greatly  by  the  hydration,  but  their  velocities  would  be. 
The  hydrated  ions,  would,  of  course,  move  more  slowly  than  the 
unhydrated. 

The  effect  of  hydration  would  obviously  be  more  pronounced  on 
freezing-point  lowering,  which  is  proportional  to  the  amount  of  solvent 
present,  than  on  conductivity.  The  following  results  taken  from  the 
work  of  Pearce1  will  show  that  this  conclusion  is  justified: 

TABLE  50. — Dissociation  from  freezing-point  lowering  and  from  conductivity. 


Salt. 

Concen- 
tration. 

Dissocia- 
tion from 
freezing- 
point 
lowering. 

Dissocia- 
tion from 
conduc- 
tivity. 

Salt. 

Concen- 
tration. 

Dissocia- 
tion from 
freezing- 
point 
lowering. 

Dissocia- 
tion from 
conduc- 
tivity. 

CaCl2  

0.01 

89.67 

CoCl2 

0.05 

90.28 

84.80 

0.05 

85.00 

80.62 

0.10 

87.36 

78.85 

0.10 

80.41 

74.35 

Mg(N03)2  

0.02 

94.90 

85.12 

SrCl2  

0.01 

91.87 

89.37 

0.05 

84.24 

78.80 

0.05 

82.65 

78.10 

0.10 

81.95 

74.78 

0.10 

81.46 

74.17 

Ba(N03)2  

0.01 

99.06 

86.37 

BaCl2  

0.01 

97.10 

90.90 

0.05 

75.18 

70.47 

0.05 

90.75 

79.80 

0.10 

62.95 

61.36 

MgCl2  

0.01 

97.10 

90.90 

Co(N03)2  

0.01 

98.65 

92.40 

0.05 

90.75 

79.78 

0.05 

88.26 

81.73 

0.10 

87.68 

73.61 

0.10 

85.48 

76.48 

SrCl2 

0.01 

91.87 

89.37 

Ni(NO3)2  

0.01 

98.03 

91.10 

0.05 

82.65 

78.08 

0.05 

83.72 

79.83 

0.10 

81.46 

74.17 

0.075 

81.32 

76.57 

Concentrated  solutions  were  also  studied  in  the  above  work;  but 
on  account  of  very  large  hydration  it  was  impossible  to  calculate  dis- 
sociation from  the  freezing-point  results. 

An  examination  of  the  above  table  will  show  that  the  dissociation 
of  dilute  solutions,  as  measured  by  the  freezing-point  method,  is  uni- 
formly greater  than  as  measured  by  the  conductivity  method.  This 
seems  to  admit  of  reasonable  explanation  in  terms  of  hydration  in 
aqueous  solution. 


'Carnegie  Inst.  Wash.,  Pub.  180;  Amer.  Chem.  Journ.,  39,  313  (1908). 


DISCUSSION   OF   EVIDENCE.  161 

EFFECT  OF  ONE  SALT  WITH  HYDRATING  POWER  ON  THE  HYDRATES  FORMED 
BY  A  SECOND  SALT  IN  THE  SAME  SOLUTION. 

The  effect  of  adding  a  salt  with  strong  hydrating  power  to  a  solution 
of  another  strongly  hydrated  salt  was  worked  out  by  Jones  and  Stine.1 
The  effect  of  adding  a  salt  with  small  hydrating  power  was  also  investi- 
gated. The  following  pairs  of  salts  were  studied :  Calcium  chloride  and 
potassium  chloride;  magnesium  chloride  and  calcium  chloride;  stron- 
tium chloride  and  calcium  chloride;  strontium  nitrate  and  magnesium 
nitrate;  calcium  nitrate  and  magnesium  nitrate;  aluminium  chloride 
and  ferric  chloride;  calcium  chloride  and  calcium  nitrate;  lithium 
bromide  and  sodium  bromide,  and  ammonium  chloride  and  potassium 
chloride,  as  examples  of  only  slightly  hydrated  salts. 

A  large  variety  of  types  of  salts  was  used.  The  first  pah-  contains  a 
binary  and  a  ternary  electrolyte  with  a  common  anion;  the  one  (calcium 
chloride)  strongly  hydrated,  the  other  (potassium  chloride)  only  slightly 
hydrated.  The  next  four  pairs  are  all  ternary  electrolytes  and  are  all 
strongly  hydrated  salts. 

Aluminium  chloride  and  ferric  chloride  are  quaternary  electrolytes 
and  strongly  hydrated.  The  two  calcium  salts  contain  a  common 
cation  and  both  are  strongly  hydrated.  The  two  bromides  contain  a 
common  anion  and  are  not  strongly  hydrated,  while  the  chlorides  of 
ammonium  and  potassium  are  very  weakly  hydrated  compounds. 

The  problem  was  obviously  a  complicated  one.  It  was  not  a 
simple  matter  to  calculate  the  composition  of  the  hydrates  formed  by 
any  one  substance  when  present  alone  in  the  solution.  It  became  far 
more  complex  and  difficult  to  calculate  the  composition  of  the  hydrates 
formed  when  two  hydrating  substances  were  present  simultaneously 
in  the  solution.  We  believe,  however,  that  this  problem  was  solved 
at  least  approximately.  It  was  found  that  the  amount  of  combined 
water  increases  with  increase  in  concentration  in  the  mixed,  as  in  the 
separate  solutions;  the  total  amount  combined  with  the  calcium 
chloride  being  less  when  the  potassium  chloride  was  present.  The 
difference  between  the  amount  of  water  combined  with  the  calcium 
chloride  when  alone  and  when  potassium  chloride  is  present  increases 
with  the  concentrations  of  the  two  constituents  of  the  mixture. 

We  found  in  general  that  when  two  hydrated  salts  were  mixed,  each 
dehydrated  the  other  to  an  amount  that  seemed  to  be  controlled  by 
mass  action.  In  order  that  this  law  should  hold,  it  would  seem  that 
the  calculated  composition  of  the  hydrates  formed  by  the  individual 
substances  must  be  approximately  correct. 

It  was  early  found  that  both  ions  and  molecules  can  form  hydrates. 
That  the  molecules  of  certain  substances  have  hydrating  power,  was 
shown  by  the  fact  that  certain  non-electrolytes  undoubtedly  combine 
with  water  in  aqueous  solution.  Ions  were,  as  a  class,  found  to  have 

2Amer.  Chem.  Journ.,  39,  313  (1908). 


162  DISCUSSION   OF   EVIDENCE. 

much  greater  hydrating  power  than  molecules.  This  conclusion  was 
confirmed  by  the  work  of  Stine.  It  was  also  shown  that  molecules 
in  aqueous  solution  can  combine  with  water,  and  in  special  cases 
molecules  may  even  have  greater  hydrating  power  than  some  ions. 

The  work  with  the  slightly  hydrated  potassium  and  ammonium 
chlorides  brought  out  a  significant  fact.  These  were  chosen,  not  to 
study  the  effect  of  one  hydrated  salt  on  the  hydration  of  another 
hydrated  salt,  but  to  study  the  effect  of  change  in  temperature  on  the 
conductivities  of  separate  solutions  of  electrolytes  and  upon  mixtures 
of  these  solutions.  For  this  purpose  it  was  necessary  to  select  a  pair  of 
salts  with  small  hydrating  power  and  also  which  do  not  form  double 
salts  with  one  another. 

If  suppression  of  ionization  were  the  only  cause  of  the  diminution 
in  conductivity  on  mixing  solutions  of  salts,  such  as  the  above,  which 
have  a  common  ion,  then  we  should  find  the  greatest  dimunition  where 
the  dissociation  is  greatest.  Since  dissociation  is  slightly  greater  at 
0°  than  at  12°,  and  slightly  greater  at  12°  than  at  25°,  we  should  expect 
to  find  greater  diminution  in  the  conductivity  at  0°  than  at  12°,  and 
greater  at  12°  than  at  25°.  Exactly  the  reverse  is  true. 

Again,  as  the  difference  in  dissociation  between  0°  to  12°  is  but  little 
greater  than  between  12°  and  25°,  we  should  expect  to  find  that  between 
these  two  ranges  of  temperature  the  driving  back  of  the  conductivity 
would  be  of  the  same  order  of  magnitude,  yet  such  is  not  the  case. 

Furthermore,  some  of  the  solutions  which  we  mixed  are  nearly  iso- 
hydric,  and  such  solutions  do  not  drive  back  each  other's  dissociation. 
Driving  back  the  dissociation  of  a  salt  by  the  addition  of  a  second  salt 
with  a  common  ion  is,  therefore,  not  the  only  cause  of  the  diminution 
in  conductivity  which  results  when  salts  with  a  common  ion  are  mixed. 

It  was  pointed  out  that  three  other  factors  may  come  into  play: 
(1)  Change  in  hydration  giving  rise  to  change  in  the  size  and  mass  of 
the  ion,  which  probably  plays  a  very  insignificant  role  in  the  above- 
named  case,  since  the  chlorides  of  ammonium  and  potassium  are  only 
slightly  hydrated.  (2)  Change  in  the  number  of  the  dissolved  parts, 
which,  however,  is  not  very  large  for  small  changes  in  temperature. 
(3)  Change  in  the  viscosity  of  the  solution  with  change  in  temperature, 
which  is  undoubtedly  a  very  prominent  factor,  hitherto  either  over- 
looked or  not  given  sufficient  prominence  in  dealing  with  the  phe- 
nomenon in  question. 

INVESTIGATIONS  IN  MIXED  SOLVENTS. 

The  study  of  the  conductivities  and  dissociations  in  pure  solvents 
was  extended  here  to  mixed  solvents.  This  phase  of  the  work  has  now 
been  in  progress  continuously  for  a  dozen  years,  and  the  results  have 
been  published  in  monographs  Nos.  80  and  180  of  the  Carnegie  Institu- 
tion of  Washington. 


DISCUSSION   OF   EVIDENCE.  163 

The  first  investigation  was  carried  out  by  Lindsay.1  He  worked 
in  water,  in  methyl,  ethyl,  and  propyl  alcohols,  and  in  mixtures  of 
these  solvents  with  one  another.  He  found,  in  certain  mixtures  of  the 
alcohols  with  water,  that  the  conductivity  of  the  dissolved  salt  was 
less  than  in  the  pure  alcohol.  The  conductivity  curves  in  mixtures 
of  methyl  alcohol  and  water  passed  through  well-defined  minima,  and 
a  conductivity  minimum  was  also  frequently  found  in  mixtures  of 
ethyl  alcohol  and  water. 

A  possible  explanation  of  the  results  in  mixtures  of  the  alcohols  with 
water  is  that  each  solvent  diminishes  the  association  of  the  other. 
Since  the  dissociating  power  of  a  solvent  is  in  general  greater  the  larger 
its  own  association,  it  follows  that  whatever  would  decrease  the  asso- 
ciation of  a  liquid  would  decrease  its  power  to  dissociate  electrolytes 
dissolved  in  it.  The  question  is,  does  one  associated  liquid  diminish 
the  association  of  another  associated  liquid? 

An  associated  liquid  tears  down  the  molecules  of  an  electrolyte 
dissolved  in  it,  into  simpler  parts  or  ions;  and  it  might  be  expected 
that  such  a  liquid  would  tear  down  the  molecules  of  another  associated 
liquid,  a  non-electrolyte,  not  into  charged  parts  or  ions,  but  into 
simpler  molecules.  The  alcohol  and  water  are  associated  liquids,  as 
has  been  shown  by  the  surface-tension  method  of  Ramsay  and  Shields.2 
Do  these  diminish  the  association  of  one  another? 

That  this  is  the  case  was  shown  by  Murray.3  He  worked  with  the 
associated  liquids,  water,  formic  acid,  and  acetic  acid.  He  determined 
the  molecular  weight  of  each  of  these  liquids  in  the  other  two,  and 
found  that  their  molecular  weights  became  smaller  the  more  dilute 
the  solutions.  This  showed  that  the  solvent,  i.  e.,  the  liquid  present 
in  the  larger  quantity,  was  tearing  down  the  molecular  complexes  of 
the  dissolved  liquid  or  the  one  present  in  smaller  quantity. 

That  the  diminution  in  the  association  of  one  associated  liquid  by 
another  associated  liquid  was  true,  was  shown  for  the  above-named 
substances  and  made  highly  probable  for  others. 

That  this  was  not  the  entire  explanation  of  the  nature  of  the  con- 
ductivity curves  in  mixtures  of  certain  alcohols  with  water,  was  brought 
out  by  the  next  investigation  in  this  field,  carried  out  here  by  Carroll.4 
He  compared  the  conductivity  curves  of  electrolytes  dissolved  in 
these  solvents,  with  the  fluidity  curves  of  the  mixtures  of  the  two  liquids 
in  question,  and  found  that  the  two  sets  of  curves  were  very  similar. 
The  minima  in  the  two  cases  occurred  in  the  same  mixture  of  the  two 
liquids.  A  careful  comparison  of  the  two  sets  of  phenomena  led  us 
to  conclude  that  the  conductivity  minima  are  largely  due  to  the  dimin- 
ished fluidity  which  takes  place  on  mixing  the  two  solvents.  The 
diminished  fluidity,  or  increased  viscosity,  would  cause  the  ions  to 
move  more  slowly,  and  hence  decrease  the  conductivity. 

1Amer.  Chem.  Journ.,  28,  329  (1902).  3Amer.  Chem.  Journ.,  30,  193  (1903). 

2Zeit.  phys.  Chem.,  12,  433  (1893).  4Ibid.,  32,  521  (1904). 


164  DISCUSSION   OF   EVIDENCE. 

At  the  end  of  the  work  done  by  Carroll,  we  seemed  justified  in 
concluding  that  the  conductivities  of  binary  electrolytes  in  such  sol- 
vents as  those  already  considered,  are  inversely  proportional  to  the 
coefficients  of  viscosity  of  the  solvent  and  are  directly  proportional  to 
the  association  of  the  solvent.  Bassett1  showed  that  silver  nitrate  in 
mixtures  of  methyl  alcohol  and  water  gave  a  conductivity  minimum 
at  both  0°  and  25°  ;  also  that  the  effect  of  one  solvent  on  the  other  was 
greater  at  0°  than  at  25°.  This  would  be  expected,  since  the  dissocia- 
tion diminishes  with  rise  in  temperature,  and  each  solvent  would 
probably  diminish  the  association  of  the  other  less,  the  smaller  its  own 
association  or  the  higher  its  temperature. 

Bingham2  not  only  measured  the  conductivities,  but  also  the  viscos- 
ities of  a  number  of  solvents  and  solutions  in  these  solvents.  He  found 
minima  in  the  conductivity  curves  in  mixtures  of  acetone  and  water. 
The  distinctly  new  feature  brought  out  by  the  work  of  Bingham  was 
that  lithium  and  calcium  nitrates  in  mixtures  of  acetone  with  methyl 
and  ethyl  alcohols  showed  a  pronounced  maximum  in  the  conduc- 
tivity curves.  This  must  be  due  either  to  an  increase  in  dissociation 
in  such  mixtures,  increasing  the  number  of  ions  present,  and  conse- 
quently increasing  the  conductivity,  or  it  must  be  due  to  a  diminution 
in  the  complexity  of  the  solvates  around  the  ions,  increasing  their 
velocities.  The  dissociation  was  measured  in  the  mixtures  in  question 
and  found  not  to  account  for  the  phenomenon.  This  eliminates 
increase  in  dissociation  and  leaves  the  other  alternative,  diminution 
in  the  complexity  of  the  solvate,  to  account  for  the  phenomenon. 

The  ion  must  drag  with  it  through  the  solvent  any  molecules  of  the 
liquid  with  which  it  had  combined.  This  would  increase  the  effective 
mass  and  diminish  its  velocity.  Anything  which  would  diminish  the 
complexity  of  the  solvate  about  the  ion  would  increase  its  velocity, 
and  consequently  the  conductivity.  We  must  therefore  conclude  that 
the  solvates  in  those  mixtures  of  acetone  with  the  alcohols  are  simplest 
where  the  conductivity  is  the  greatest. 

Rouiller3  studied  both  the  velocities  of  the  ions  and  the  conductivities 
of  electrolytes  in  mixtures  of  acetone  with  the  alcohols.  Silver  nitrate 
in  methyl  alcohol  and  acetone  gave  a  decided  maximum  of  conduc- 
tivity. His  work  on  the  velocities  of  the  ions  in  these  mixtures  indi- 
cated that  the  above  explanation  of  the  maxima  offered  by  Jones 
and  Bingham  was  correct;  there  is  a  change  in  the  complexity  of  the 
solvate  about  the  ion. 

McMaster4  extended  the  work  in  the  same  solvents  used  by  Bingham  — 
water,  methyl  alcohol,  ethyl  alcohol,  and  acetone  —  and  in  mixtures 
of  these  with  one  another.  He  found  conductivity  results  of  the  same 
general  character  as  those  obtained  by  the  earlier  workers.  Conduc- 


r.  Chem.  Journ.,  32,  409  (1904).  3Ibid.,  36,  443  (1906). 

*Ibid.,  34,  481  (1905).  *Ibid.,  326  (1906). 


DISCUSSION   OF   EVIDENCE.  165 

tivity  minima  were  found  in  mixtures  of  the  alcohols  with  water  and 
acetone  with  water.  Conductivity  maxima  were  obtained  with  lithium 
bromide  in  mixtures  of  methyl  or  ethyl  alcohol  with  acetone.  Cobalt 
chloride  in  mixtures  of  acetone  with  ethyl  alcohol  also  showed  a  maxi- 
mum. Jones  and  McMaster  reached  the  same  conclusion  from  their 
work  that  had  been  reached  by  Jones  and  Bingham.  Change  in  the 
complexity  of  the  solvate  formed  by  the  ion  in  different  mixtures  of 
solvents  is  an  important  factor  in  determining  the  conductivity  maxima. 

A  point  of  interest  brought  out  by  the  work  of  McMaster  was  in 
connection  with  the  temperature  coefficients  of  conductivity  in  non- 
aqueous  solutions.  The  bearing  of  temperature  coefficients  of  con- 
ductivity on  the  solvate  theory  of  solution  has  already  been  discussed. 
With  rise  in  temperature  the  hydrates  about  the  ions  became  simpler. 
The  mass  and  probably  the  size  of  the  ion  thus  became  less,  and  it  moves 
faster  the  higher  the  temperature,  thus  increasing  the  conductivity. 

McMaster  found  that  cobalt  chloride  in  certain  mixtures  of  acetone 
with  the  alcohols  showed,  at  ordinary  temperatures,  negative  temper- 
ature coefficients  of  conductivity.  What  does  this  mean?  The  solvent 
becomes  less  viscous  with  rise  in  temperature,  thus  increasing  the 
velocity  of  the  ions;  and  the  solvates  become  simpler,  which  also 
increases  the  velocity  with  which  the  ions  move. 

With  rise  in  temperature,  on  the  other  hand,  the  association  of  the 
solvent,  and  consequently  its  dissociating  power,  becomes  less. 

The  above  two  influences  work  counter  to  one  another.  Negative 
temperature  coefficients  of  conductivity  mean  that  the  latter  influence 
overcomes  the  former.  The  alcohols  used  and  acetone  are  highly 
associated  liquids.  Rise  in  temperature  diminishes  their  association 
and  consequently  their  dissociating  power. 

A  solution  of  cobalt  chloride  in  a  75  per  cent  mixture  of  acetone  with 
methyl  alcohol,  the  solution  being  ^j-  normal,  had  a  zero  temperature 
coefficient  of  conductivity. 

A  number  of  points  of  interest  were  brought  out  by  the  next  investi- 
gator, Veazey.1  He  worked  with  solutions  of  salts  in  water,  methyl 
alcohol,  ethyl  alcohol,  acetone,  and  in  binary  mixtures  of  these  solvents 
with  one  another.  The  minimum  in  conductivity  was  found  to  be  a 
more  general  phenomenon  than  had  been  supposed  from  the  earlier 
work.  It  had  long  been  known  that  mixtures  of  methyl  alcohol  and 
water  or  ethyl  alcohol  and  water,  are  more  viscous  than  either  of  the 
pure  solvents  alone.  A  rational  explanation  of  this  phenomenon  was 
suggested — alcohol  and  water  are  strongly  associated  liquids.  When 
two  associated  liquids  are  mixed  each  diminishes  the  association  of 
the  other.  The  larger  molecules  are  thus  broken  down  into  smaller 
molecules,  which  increases  the  frictional  surfaces  when  these  molecules 
move  over  one  another  as  they  do  in  measuring  viscosity.  The  result 

'Amer.  Chem.  Journ.,  37,  405  (1907).     Zeit.  phys.  Chem.,  61,  641  (1908);  62.  44  (1908). 


166  DISCUSSION   OF   EVIDENCE. 

would  be  to  increase  the  viscosity  of  the  mixture  over  that  of  either 
pure  solvent. 

Maxima  in  the  conductivity  of  electrolytes  in  the  mixed  solvents 
were  shown  to  correspond  to  maxima  in  the  fluidity  of  the  mixed 
solvents.  Maxima  in  fluidity  are  probably  due  to  an  increase  in  the 
size  of  the  molecules  of  the  solvent,  due  to  a  combination  of  one 
solvent  with  the  other.  This  would  diminish  the  viscosity  and  conse- 
quently increase  the  velocity  of  the  ions,  which  would  increase  the 
conductivity.  This  factor  must  also  be  taken  into  account  in  explain- 
ing conductivity  maxima. 

The  temperature  coefficients  of  conductivity  in  the  above-named 
mixtures  of  liquids  with  water  are  a  maximum  in  the  25  and  50  per 
cent  mixtures.  These  are  just  about  the  mixtures  in  which  the  sol- 
vents have  the  least  association.  The  molecules  of  the  solvents  being 
in  the  simplest  condition,  would  be  most  favorable  for  chemical  action. 
In  such  mixtures  the  solvents  probably  combine  to  the  greatest  extent 
with  the  dissolved  substance — the  solvation  is  at  a  maximum.  The 
effect  of  rise  in  temperature  breaking  down  these  solvates  would  there- 
fore be  a  maximum  where  solvation  is  a  maximum.  Solutions  of 
potassium  sulphocyanate  have  greater  conductivity  in  acetone  than 
in  water.  This  was  shown  to  be  due  to  the  greater  fluidity  of  the 
acetone. 

This  same  salt  when  dissolved  in  water  lowers  the  viscosity  of  the 
water.  An  examination  of  the  literature  showed  that  certain  salts  of 
potassium  and  salts  of  rubidium  and  caesium  are  practically  the  only 
ones  known  to  lower  the  viscosity  of  water.  In  the  case  of  certain 
salts  of  potassium  the  positive  effect  of  the  anion  on  the  viscosity  of 
water  may  more  than  offset  the  negative  effect  of  the  potassium  ion. 

The  following  explanation  of  the  above-named  phenomenon  was 
suggested.  If  the  atomic  volume  of  the  ions  dissolved  in  the  solvent 
was  larger  than  the  molecular  volume  of  the  solvent,  the  larger  ions 
would  diminish  the  size  of  the  frictional  surfaces  coming  in  contact 
and  would  lower  the  viscosity. 

It  is  wTell  known  that  potassium,  rubidium,  and  csesium  occupy  the 
maxima  on  the  atomic-volume  curve,  and  have  much  larger  atomic 
volumes  than  any  other  known  elements.  Potassium  has  a  smaller 
atomic  volume  than  rubidium,  and  rubidium  than  caesium.  Potassium 
chloride  lowers  the  viscosity  of  water  less  than  rubidium  chloride,  and 
rubidium  chloride  less  than  caesium  chloride. 

If  we  study  the  salts  which  raise  the  viscosity  of  water,  we  will  find, 
in  general,  that  the  amount  of  increase  in  the  viscosity  bears  a  relation 
to  the  atomic  or  ionic  volumes  of  the  dissolved  substances.  Smaller 
ions  tend  to  increase  the  viscosity  of  water  more  than  larger  ones.  It 
would  therefore  seem  that  the  above  explanation  contains  a  large 
element  of  truth. 


DISCUSSION   OF   EVIDENCE.  167 

The  work  already  discussed  in  mixed  solvents  is  all  recorded  in 
Publication  No.  80  of  the  Carnegie  Institution  of  Washington.  The 
results  of  the  following  five  investigations  are  recorded  in  Publication 
No.  180  of  the  Carnegie  Institution  of  Washington. 

The  problem  of  measuring  dissociation  in  non-aqueous  solvents  is 
a  difficult  one.  The  freezing-point  method  is  frequently  not  applicable. 
Many  common  solvents,  such  as  the  alcohols,  freeze  at  temperatures 
which  are  too  widely  removed  from  the  ordinary  temperature  of  the 
laboratory  to  measure  with  sufficient  accuracy.  The  boiling-point 
method  could  be  used  only  with  fairly  concentrated  solutions.  Dilute 
solutions  produce  such  a  slight  rise  in  the  boiling-point  that  this  small 
quantity  can  not  be  measured  with  a  very  high  degree  of  accuracy. 
The  boiling-point  method  has  the  further  disadvantage  of  being  so 
largely  affected  by  slight  changes  in  the  barometer. 

The  hope  of  measuring  conductivity  in  non-aqueous  solvents  in 
general  seemed  to  rest  in  the  conductivity  method.  This  method  as 
ordinarily  applied  would  not  be  satisfactory.  The  dilution  at  which 
complete  dissociation  would  be  reached  in  such  solvents  is  so  great 
that  the  Kohlrausch  method  in  any  such  form  as  he  left  it  could  not 
be  applied  to  the  problem. 

The  conductivity  method  was  greatly  improved  by  Kreider;1  the 
greatest  improvement  being  in  the  form  of  cell  employed.  With  the 
improved  method  Kreider  studied  the  dissociations  of  a  number  of 
salts  in  methyl  and  ethyl  alcohols  and  in  mixtures  of  these  solvents 
with  water.  He  measured  the  conductivities  of  solutions  as  dilute 
as  100,000  liters. 

uoo  methyl  alcohol 
He  found  the  following  relation:  77— j — j — ITT~=  constant. 

When  a  salt  is  dissociated  in  each  of  two  solvents,  for  the  same  con- 
centration of  the  salt  there  are  the  same  number  of  ions  in  the  two 
solutions.  Conductivity  is  a  function  of  the  number  of  the  ions  and 
their  velocities.  When  numbers  of  the  ions  are  constant,  as  in  this 
case,  conductivity  is  a  function  of  the  relative  velocities  of  the  ions. 
The  velocity  of  an  ion  is  conditioned  by  its  mass  and  volume  and  by 
the  fluidity  of  the  solvent.  If  the  masses  and  volumes  of  the  ions  in 
the  two  solvents  are  constant,  the  velocities  of  the  ions  should  vary 
as  the  fluidities  of  the  solvents.  The  ratio  between  the  values  of  ^» 
in  the  two  solvents  should  be  the  same  as  the  ratio  between  the  fluidities 
of  these  solvents.  This  was,  however,  found  not  to  be  the  case.  The 
bearing  of  this  fact  on  the  condition  of  the  ions  in  the  two  solvents  in 
question  is  important.  This  shows  that  the  mass  and  probably  the 
volume  of  the  solvated  ion  must  differ  in  the  two  solvents. 

The  ratio  between  the  values  of  //«,  for  a  salt  in  the  two  solvents, 
compared  with  the  ratio  between  the  fluidities  of  the  two  solvents, 

'Amer.  Chem.  Journ.,  45,  282  (1911). 


168  DISCUSSION   OF   EVIDENCE. 

would  give  an  approximate  idea  of  the  relative  solvation  of  the  ions 
in  the  two  solvents  in  question. 

This  method  will  be  still  further  applied  to  the  problem  of  solvation 
in  non-aqueous  solvents. 

Mahin1  studied  electrolytes  in  ternary  mixtures  of  the  alcohol  with 
water,  and  obtained  results  of  the  same  general  character  as  those 
found  in  binary  mixtures  of  these  solvents.  He  then  took  up  work  in 
binary  mixtures,  one  constituent  being  acetone.  Acetone  was  studied 
primarily  because  it  is  an  exceptional  solvent  in  many  of  its  properties. 
Substances  dissolved  in  acetone  are  largely  polymerized,  and  acetone 
has  at  the  same  time  considerable  dissociating  power.  Furthermore, 
acetone  is  a  solvent  with  small  viscosity,  and  it  was  desired  to  see 
whether  the  relations  found  for  solvents  with  larger  viscosity  would 
hold  here.  The  curve  for  conductivity  and  for  fluidity  were  worked 
out  and  the  two  compared. 

It  was  found  that  the  product  of  molecular  conductivity  and  vis- 
cosity is  nearly  a  constant  at  complete  dissociation.  This  means  that 
for  completely  dissociated  solutions  in  acetone  the  curves  of  molecular 
conductivity  are  similar  to  those  of  fluidity — conductivity  being  in- 
versely proportional  to  viscosity.  This  relation  is  of  interest  in  that 
it  holds  in  a  solvent  with  such  small  viscosity  as  acetone. 

Relations  such  as  those  referred  to  above  having  been  found  to  hold 
in  a  solvent  with  such  small  viscosity  as  acetone,  the  question  arose, 
do  such  relations  obtain  in  a  highly  viscous  solvent  like  glycerol? 
Glycerol  not  only  has  a  very  high  viscosity,  but  is  an  excellent  solvent, 
and  has  a  large  dielectric  constant,  which  means  that  it  has  consider- 
able dissociating  power.  Glycerol  is  fairly  strongly  associated,  which 
also  indicates  considerable  dissociating  power. 

The  first  investigation  in  glycerol  as  a  solvent  was  carried  out  by 
Schmidt.2  He  measured  the  conductivities  of  solutions  of  certain 
salts  in  glycerol,  and  in  mixtures  of  glycerol  with  water  and  with  methyl 
and  ethyl  alcohols.  The  conductivities  were  measured  at  different 
temperatures.  The  most  striking  relation  noted  was  the  enormous 
magnitude  of  the  temperature  coefficients  of  conductivity  of  electro- 
lytes dissolved  in  glycerol.  This  was  shown  to  be  due  to  the  rapid 
decrease  in  the  viscosity  of  glycerol  with  rise  in  temperature. 

It  was  shown  that  where  glycerol  is  mixed  with  water  or  the  alcohols, 
there  is  a  breaking  down  of  the  association  of  each  solvent  by  the  other, 
and  a  consequent  diminution  in  the  dissociating  power.  Solutions  of 
potassium  iodide  in  25  and  50  per  cent  mixtures  of  glycerol  and  water 
lowered  the  viscosity  of  these  solvents.  This  salt  does  not  lower  the 
viscosity  of  glycerol,  but  of  the  mixtures.  The  meaning  of  negative 
viscosity  effects  was  discussed  in  the  work  of  Veazey.  While  Schmidt 
did  not  study  any  salt  which  lowers  the  viscosity  of  pure  glycerol,  he 

iAmer.  Chem.  Journ.,  41,  433  (1909) ;  Zeit.  phys.  Chem.,  69,  389  (1909). 
2Amer.  Chem.  Journ.,  42,  37  (1909). 


DISCUSSION   OF   EVIDENCE.  169 

found  that  the  effect  of  the  salt  on  the  viscosity  of  pure  glycerol  was 
inversely  as  the  molecular  volume  or  atomic  volumes  of  the  constitu- 
ents of  the  salt.  This  was  in  keeping  with  the  explanation  offered  by 
Jones  and  Veazey  to  account  for  the  changes  in  the  viscosity  of  the 
solvent  by  the  dissolved  substance.  A  comparison  of  the  conduc- 
tivity and  fluidity  curves  shows  that  the  two  run  nearly  parallel. 
Although  glycerol  has  about  1,000  times  the  viscosity  of  methyl  alcohol, 
yet,  from  the  work  of  Schmidt,  the  same  general  relations  obtain  here 
that  hold  for  the  far  less  viscous  solvents. 

The  work  of  Schmidt  was  continued  by  Guy.1  He  worked  with  a 
much  larger  number  of  salts,  and  over  the  temperature  range  25°  to 
75°.  He  studied  not  only  solutions  in  glycerol,  but  in  mixtures  of 
glycerol  with  water,  with  methyl,  and  with  ethyl  alcohols. 

Guy  found  also  enormous  temperature  coefficients  of  conductivity. 
This  may  be  due  to  either  of  two  causes :  a  change  in  dissociation  with 
rise  in  temperature,  or  a  change  in  the  velocity  of  the  ions.  We  know 
the  order  of  magnitude  of  the  change  in  dissociation  with  rise  in  tem- 
perature, and  it  is  small.  The  chief  cause  of  the  large  temperature 
coefficients  of  conductivity  in  glycerol  is,  then,  an  increase  in  the 
velocities  with  which  the  ions  move.  As  we  have  seen,  this  may  be 
due  to  a  decrease  in  the  viscosity  of  the  solvent  with  rise  in  temperature, 
or  may  be  caused  by  a  breaking  down  of  complex  solvates  about  the  ions. 

While  the  viscosity  of  glycerol  increases  rapidly  with  rise  in  tem- 
perature, this  alone  would  not  account  for  the  magnitude  of  the  temper- 
ature coefficients  of  conductivity  of  glycerol  solutions.  There  seems 
to  be  good  evidence  for  the  formation  of  glycerolates  in  solutions  in 
glycerol.  The  temperature  coefficients  of  conductivity  in  glycerol  are 
greater  at  high  than  at  low  dilution.  Jones  has  pointed  out  that  this 
would  be  expected  from  the  solvate  theory.  The  more  dilute  the 
solution  the  more  complex  the  solvate;  the  more  complex  the  solvate 
the  greater  the  change  in  its  complexity  with  rise  in  temperature. 

Further,  salts  of  calcium,  strontium,  and  barium  have  larger  tem- 
perature coefficients  of  conductivity  than  those  of  sodium,  potassium, 
and  ammonium.  The  former  are  strongly  hydrated,  the  latter  weakly 
hydrated  substances.  It  would  seem  that  the  former  are  more  strongly 
glycerolated  than  the  latter.  Salts  which  have  approximately  the 
same  hydrating  power  have  temperature  coefficients  of  conductivity 
in  glycerol  of  the  same  order  of  magnitude,  indicating  the  same  order 
of  magnitude  of  glycerolation.  Work  in  the  mixed  solvents  indicates 
that  water  diminishes  the  association  of  glycerol. 

Solutions  of  salts  in  glycerol  have  in  general  greater  viscosity  than 
pure  glycerol.  Guy,  however,  found  marked  exceptions  to  this  rela- 
tion. Salts  of  rubidium  lowered  the  viscosity  of  glycerol.  Ammonium 
bromide  and  iodide  also  lowered  the  viscosity  of  this  solvent.  That 

.  Chem.  Journ.,  46,  131  (1911). 


170  DISCUSSION   OF   EVIDENCE. 

rubidium  should  lower  the  viscosity  of  glycerol  is  in  keeping  with 
what  was  found  in  aqueous  solutions.  Salts  of  rubidium  and  caesium 
and  some  salts  of  potassium  lowered  the  viscosity  of  water.  This  has 
already  been  explained  as  due  to  the  large  atomic  volumes  of  these 
elements.  The  same  explanation  holds  for  solutions  in  glycerol. 

Davis1  continued  the  work  of  Guy,  studying  especially  the  effect 
of  salts  on  the  viscosity  of  glycerol.  He  repeated  the  work  with 
ammonium  iodide  and  obtained  the  same  result  that  had  been  earlier 
found  by  Guy.  He  studied  rubidium  chloride,  bromide,  iodide,  and 
nitrate,  and  showed  that  these  lowered  the  viscosity  of  glycerol.  The 
rubidium  salts  lower  the  viscosity  of  glycerol  to  such  an  extent  that 
they  appreciably  increase  their  own  conductivity  in  this  solvent. 

Comparing  the  effects  of  the  chloride,  bromide,  and  iodide  of  rubid- 
ium on  the  viscosity  of  glycerol,  Davis  found  that  the  chloride  has  the 
least  effect,  the  bromide  next,  the  iodide  the  greatest.  He  showed  that 
this  was  in  the  same  order  as  the  molecular  volumes  of  the  salts  in 
question.  The  results  obtained  with  glycerol  were,  then,  analogous  to 
those  obtained  with  water,  both  with  respect  to  viscosity  and  solution. 

SPECTROSCOPIC  EVIDENCE  BEARING  ON  THE  SOLVATE  THEORY  OF  SOLUTION. 
WORK  OF  JONES  AND  UHLER. 

Work  on  the  absorption  spectra  of  solutions  has  now  been  in  progress 
in  my  laboratory  continuously  for  eight  years.  This  work  was  under- 
taken in  connection  with  its  bearing  on  the  solvate  theory  of  solution. 
What  connection  is  there  between  solvation  and  the  power  of  solu- 
tions to  absorb  light? 

It  is  well  known  that  absorption  of  light  means  that  the  wave- 
lengths of  light  set  something  vibrating  with  periods  the  same  as  their 
own.  Selective  absorption  of  light  or  the  absorption  of  certain  wave- 
lengths of  light  means  that  the  wave-lengths  absorbed  set  something 
vibrating  with  their  own  periods.  Absorption  of  light  is,  then,  a 
resonance  phenomenon.  Absorption  of  light  by  a  dissolved  substance 
means  that  something  in  the  solution  must  be  thrown  into  resonance 
with  the  light — must  be  set  vibrating  with  the  same  periods  as  the 
light-waves.  Many  dissolved  substances  absorb  only  certain  wave- 
lengths. This  means  that  those  particular  wave-lengths  of  light  find 
something  in  the  solution  which  they  can  set  vibrating  with  their  own 
periods.  Transparency  means  lack  of  resonance,  opacity  means  reso- 
nance. The  color  of  any  given  solution  is  determined  by  the  wave- 
lengths of  light  which  are  not  absorbed.  A  red  solution  is  one  which 
allows  the  long  wave-lengths  to  pass  through.  A  blue  solution  is  one 
which  allows  the  short  wave-lengths  to  pass  through.  That  particle 
in  solution  which  is  thrown  into  resonance  by  the  light  is  called  the 

*Zeit.  phys.  Chem.,  81,  68  (1912). 


DISCUSSION   OF   EVIDENCE.  171 

resonator.  This  was  formerly  supposed  to  be  the  molecule  or  the  ion, 
but  is  now  thought  to  be  the  electron.  Whatever  the  nature  of  the 
resonator,  the  absorption  of  light  by  dissolved  substance  is  due  to  it. 

The  line  of  thought  which  led  us  to  take  up  the  study  of  the  absorp- 
tion spectra  of  solutions  in  connection  with  the  solvate  theory  of 
solution  is  the  following:  The  absorption  of  light  being  due  to  a 
resonator,  this  would  have  different  resonance  when  anhydrous  than 
when  combined  with  molecules  of  the  solvent.  In  general,  the  reso- 
nance would  be  different  when  the  resonator  was  unsolvated  than 
when  it  was  solvated.  The  color  of  the  solution  being  due  to  the 
resonator,  the  solution  could  reasonably  be  expected  to  have  different 
color  when  the  resonator  was  solvated  than  when  it  was  unsolvated. 
The  study  of  the  color  of  solutions,  and  the  changes  in  the  color  when 
the  resonator  underwent  changes  in  solvation,  might  give  some  clue 
to  the  changes  in  solvation. 

It  is  a  comparatively  simple  matter  to  change  solvation  in  solution; 
it  is  only  necessary  to  change  the  concentration  of  the  solution.  The 
more  dilute  the  solution  the  more  complex  the  solvates  formed.  We  shall 
see  that  this  often  produces  a  marked  change  in  the  absorption  spectra. 
We  can  diminish  the  complexity  of  the  solvates  by  raising  the  temper- 
ature. This  also  frequently  produces  marked  changes  in  the  absorption. 
Addition  of  a  dehydrating  agent  will  change  the  hydration  of  any  given 
salt.  This  frequently  changes  the  absorption  spectra  and  the  color  of  a 
solution ;  and  there  are  many  other  ways  of  changing  solvation.  These 
frequently  produce  concomitant  changes  in  the  absorption  spectra. 

A  salt  dissolved  in  water  may  form  hydrates,  in  alcohol  alcoholates, 
in  acetone  acetonates,  in  glycerol  glycerolates,  etc.  We  should  expect 
these  different  solvates  to  affect  the  resonator  or  resonators  differently. 
We  shall  see  that  this  is  true. 

With  this  idea  in  mind,  work  was  begun  in  my  laboratory  on  the 
study  of  the  absorption  spectra  of  solutions.  The  first  investigation 
was  carried  out  by  Dr.  Uhler  and  myself.  Our  work  consisted  largely 
in  devising  a  method  and  apparatus  for  studying  the  property  of 
solutions  to  absorb  light.  The  key  to  the  method  consisted  in  using 
a  grating  instead  of  a  prism  spectroscope.  This  gave  much  greater 
dispersion,  and  brought  out  many  new  lines  and  bands.  A  form  of 
cell  was  devised  for  holding  solutions  in  non-aqueous  solvents  which 
avoided  the  use  of  all  cement.  The  details  of  this  phase  of  the  work 
are  all  given  in  Publication  No.  60  of  the  Carnegie  Institution  of 
Washington.  We  studied  the  effect  on  the  absorption  spectra  of 
increasing  the  concentration  of  the  solution,  and  found  that,  in  general, 
the  effect  of  increasing  the  concentration  of  the  solution  was  to  widen 
the  absorption  bands.  As  the  solvates  became  simpler  the  absorption 
bands  became  broader. 

Another  method  of  simplifying  the  hydrates  existing  in  an  aqueous 
solution  was  to  add  a  dehydrating  agent  in  the  form  of  a  second  salt. 


172  DISCUSSION   OF   EVIDENCE. 

It  was  found  that  this  also  produced  a  widening  of  the  absorption 
bands.  This  was  in  keeping  with  the  effect  of  increasing  the  concen- 
tration of  the  solution,  which  also  simplified  the  hydrates. 

Jones  and  Uhler  also  studied  the  effect  of  adding  water  to  solutions 
in  non-aqueous  solvents.  Thus,  water  was  added  to  solutions  in 
methyl  and  ethyl  alcohols,  acetone,  etc.  The  effect  of  adding  water 
was  to  narrow  the  absorption  bands.  All  of  these  results  were  regarded 
as  in  keeping  with  the  solvate  theory  of  solution. 

WORK  OF  JONES  AND  ANDERSON. 

The  work  of  Jones  and  Uhler  on  the  absorption  spectra  of  solutions 
was  greatly  extended  in  a  number  of  directions  by  Jones  and  Anderson.1 
They  worked  with  salts  of  cobalt,  nickel,  copper,  iron,  chromium,  neo- 
dymium,  praseodymium,  and  erbium.  Only  that  phase  of  the  work  will 
be  discussed  here  which  bears  on  the  solvate  theory  of  solution. 

We  will  first  consider  the  results  with  salts  of  cobalt.  There  is  a 
region  of  one-sided  absorption  in  the  ultra-violet.  This  band  narrows 
with  dilution,  but  remains  of  approximately  constant  width  when  the 
number  of  molecules  in  the  path  of  the  beam  of  light  is  kept  constant, 
indicating  that  the  absorbers  here  are  the  undissociated  molecules. 

The  band  X3300  disappears  rapidly  with  increase  in  the  dilution  of 
the  solution,  even  when  the  number  of  molecules  in  the  path  of  the 
beam  of  light  is  kept  constant.  This  band  increases  rapidly  in  intensity 
with  rise  in  temperature,  and  can  be  accounted  for  best  by  assuming 
that  it  is  due  to  a  relatively  simple  hydrate.  It  is  well  known  that 
rise  in  temperature  breaks  down  complex  hydrates  into  simpler  ones, 
which  would  give  rise  to  the  band;  and,  further,  increase  in  dilution 
produces  more  and  more  complex  hydrates.  These  would  cause  the 
disappearance  of  a  band  due  to  simpler  hydrates. 

The  green  cobalt  band  can  not  be  due  to  the  cobalt  ions,  since  it  is 
not  most  intense  where  the  number  of  cobalt  ions  is  the  greatest.  The 
width  of  this  band  does  not  vary,  if  the  light  is  passed  through  such 
depths  of  the  solution  that  the  product  of  the  concentration  multiplied 
by  the  depth  is  kept  constant.  This  would  indicate  that  this  band 
is  due  to  the  cobalt  atom,  whether  combined  as  in  the  molecule  or 
dissociated  as  an  ion. 

The  absorption  in  the  red  is  characteristic  of  concentrated  solutions 
alone.  This  would  show  that  it  is  not  due  to  the  cobalt  ion.  We 
might  suppose  that  it  was  due  to  aggregates  of  molecules;  but  this 
view  is  not  tenable,  since  the  absorption  in  the  red  increases  with  rise 
in  temperature,  which  breaks  down  such  aggregates.  High  tempera- 
ture and  great  concentration  favor  the  formation  of  simple  hydrates 
and  also  increase  the  absorption  in  the  red.  The  red  absorption  can 
therefore  be  accounted  for  as  due  to  simple  hydrates  in  solution. 

Carnegie  Inst.  Wash.  Pub.  No.  110;  Amer.  Chem.  Journ.,  41,  163  (1909). 


DISCUSSION   OF   EVIDENCE.  173 

Jones  and  Anderson  also  did  some  work  on  the  absorption  spectra  of 
cobalt  salts  in  certain  non-aqueous  solvents.  The  green  band  appeared 
in  all  of  the  non-aqueous  solutions  studied.  This  is  just  what  would 
be  expected  if  this  band  is  due  to  the  cobalt  atom.  The  intensity  of 
this  band  in  non-aqueous  solvents  was  found  to  be  proportional  to 
the  concentration. 

The  red  absorption  is  more  intense  in  the  non-aqueous  solvents 
than  in  water,  the  intensity  increasing  from  methyl  alcohol  to  ethyl 
alcohol  to  acetone.  With  increase  in  the  dilution  the  band  narrows 
rapidly  in  methyl  alcohol,  more  slowly  in  ethyl  alcohol,  and  remains 
nearly  constant  in  acetone.  All  of  these  facts  are  in  accord  with  the 
view  that  this  band  is  due  to  simple  solvates.  We  should  expect  the 
power  to  form  solvates  to  be  greater  for  methyl  alcohol  than  for  ethyl  alco- 
hol, and  greater  for  ethyl  alcohol  than  for  acetone.  In  ethyl  alcohol 
and  acetone  at  ordinary  temperatures  most  of  the  solvates  are  probably 
simple  enough  to  absorb  in  the  red,  while  in  methyl  alcohol  this  is 
the  case  only  at  elevated  temperatures  or  in  concentrated  solutions. 

The  nickel  absorption  bands  are  similar  in  their  behavior  to  the 
green  band  of  cobalt.  The  absorption  of  nickel  salts  seems  to  be 
largely  a  function  of  the  nickel  atom.  The  widening  of  the  band 
X  3960,  with  concentration  and  with  hydrating  agents,  indicates  that 
the  simplest  hydrates  have  a  somewhat  different  absorption  from  the 
more  complex.  The  ultra-violet  absorption  of  copper  salts  decreases 
rapidly  with  dilution,  when  we  keep  the  product  of  depth  of  layer  and 
concentration  constant.  This  would  indicate  that  this  absorption  is 
not  due  to  the  ions,  but  must  be  due  in  some  way  to  the  molecules. 
The  absorption  decreases  with  the  dilution,  even  when  the  molecules 
in  the  path  of  the  light  are  kept  constant.  This  would  indicate  that 
the  absorbing  power  of  molecules  is  affected  by  the  surroundings. 

The  increase  in  the  absorption  with  concentration  when  the  mole- 
cules are  kept  constant  might  be  due  to  the  formation  of  molecular 
aggregates  or  might  be  due  to  solvates.  To  decide  between  these  two 
alternatives  we  must  take  into  account  the  effect  of  rise  in  tempera- 
ture on  the  absorption.  Rise  in  temperature  increases  the  absorption, 
but  rise  in  temperature  breaks  down  the  molecular  aggregates.  There- 
fore, this  absorption  can  not  be  due  to  aggregates.  Solvates  become 
simpler  both  by  rise  in  temperature  and  by  increase  in  the  concentra- 
tion of  the  solution.  Both  should  produce  the  same  effect  on  the 
absorption  if  this  absorption  is  due  to  simple  solvates,  and  such  is  the 
fact.  We  must  therefore  conclude  that  the  ultra-violet  absorption  of 
copper  salts  is  due  to  simpler  hydrates. 

For  equal  concentration  the  ultra-violet  absorption  of  copper  salts 
is  least  in  the  aqueous  solutions,  and  increases  as  we  pass  from  methyl 
to  ethyl  alcohol.  Further,  the  change  in  this  absorption  with  dilution 
is  greatest  for  the  aqueous  solutions,  and  decreases  as  we  pass  to  methyl 


174  DISCUSSION   OF   EVIDENCE. 

and  ethyl  alcohols.  These  facts  are  just  what  would  be  expected  if 
this  absorption  was  due  to  simpler  solvates,  since  the  power  to  form 
solvates  is  greater  for  water  than  for  either  of  the  alcohols,  and  greater  for 
methyl  than  for  ethyl  alcohol.  For  equal  concentrations  the  solvates 
would  decrease  in  complexity  as  we  pass  from  water  to  methyl  alcohol. 
Further,  increase  in  dilution  would  change  the  complexity  of  the  sol- 
vates more  in  aqueous  solutions  than  in  solutions  in  either  of  the  alcohols. 
The  above  conclusion  is,  then,  in  perfect  accord  with  all  of  the  facts. 

The  absorption  of  copper  salts  in  the  red  narrows  when  the  product 
of  concentration  and  depth  of  absorbing  layer  is  kept  constant,  but 
widens  when  the  molecules  are  kept  constant.  Its  intensity  varies 
far  more  with  change  in  concentration  than  with  change  in  solvent. 
This  absorption  must  be  due  to  the  atom,  and  is  affected  comparatively 
slightly  by  the  surroundings  of  the  atom.  The  copper  absorption  in 
the  red  is,  then,  less  affected  by  solvates  than  the  absorption  by  copper 
in  the  ultra-violet.  The  feature  of  the  work  of  Jones  and  Anderson, 
which  bears  most  directly  on  the  solvate  theory  of  solution,  came  out 
as  the  result  of  studying  the  absorption  spectra  of  solutions  of  salts 
of  neodymium  and  praseodymium,  and  especially  of  neodymium. 

Neodymium  chloride  was  found  to  have  quite  different  absorption 
in  water  from  what  it  had  in  methyl  alcohol.  This  made  it  desirable 
to  study  the  absorption  spectrum  of  this  salt  in  mixtures  of  methyl 
alcohol  and  water.  By  changing  the  composition  of  the  mixtures  of 
the  two  solvents,  we  could  see  how  the  spectra  corresponding  to  the 
two  solvents  would  change. 

It  was  found  that  when  the  proper  mixture  of  alcohol  and  water 
was  used,  the  two  spectra  (the  one  corresponding  to  the  alcoholic 
solution  and  the  other  to  the  aqueous  solution)  coexisted  on  the  plate. 
When  the  amount  of  water  in  the  mixed  solvents  increased,  the  "water 
spectrum"  came  out  more  strongly;  when  the  amount  of  alcohol 
present  was  increased,  the  " alcohol  spectrum"  came  out  more  strongly. 
When  the  amount  of  water  present  exceeds  15  or  20  per  cent,  we  have 
only  the  "  water  spectrum."  As  the  amount  of  water  is  still  further 
decreased  by  the  addition  of  more  alcohol,  the  spectrum  consists  of 
the  "water  spectrum"  and  the  "alcohol  spectrum"  superposed.  As 
the  amount  of  water  is  diminished  below  15  per  cent,  the  intensity  of 
the  water  spectrum  becomes  less  and  less  and  the  intensity  of  the 
alcohol  spectrum  greater  and  greater. 

A  question  of  importance  in  the  present  connection  is  this :  Does  the 
"water  spectrum"  gradually  change  over  into  the  alcohol  spectrum" 
as  the  amount  of  alcohol  present  is  increased,  or  do  we  have  here  two 
separate  and  distinct  spectra,  the  one  corresponding  to  the  aqueous 
solution,  and  the  other  to  the  alcoholic? 

To  test  this  point,  we  worked  with  fairly  dilute  solutions  of  neody- 
mium chloride  in  water,  in  methyl  alcohol,  and  in  mixtures  of  water 


DISCUSSION   OF   EVIDENCE.  175 

and  methyl  alcohol.  The  object  in  using  dilute  solutions  was  to  be 
able  to  study  the  structure  of  the  bands  in  the  different  solvents.  In 
the  more  dilute  solutions  the  several  parts  of  any  given  band  would 
come  out  clearly  and  could  be  measured.  The  result  was  to  show 
that  the  " alcohol  spectrum"  was  quite  different  from  the  "water 
spectrum."  It  had  different  components  and  they  were  arranged  in 
a  different  way  within  the  bands. 

In  mixed  solvents,  then,  the  two  spectra  coexisted,  and  we  did  not 
have  the  one  passing  over  into  the  other  as  we  changed  the  composi- 
tion of  the  mixture  of  alcohol  and  water.  The  "water"  spectrum  and 
"methyl  alcohol"  spectrum  had  equal  intensities  when  the  mixture  of 
the  water  and  methyl  alcohol  contained  from  6  to  8  per  cent  of  water. 

Neodymium  nitrate  shows  change  in  the  spectra  analogous  to  those 
manifested  by  the  chloride,  when  dissolved  in  mixtures  of  water  and 
one  of  the  non-aqueous  solvents.  The  change  with  the  nitrate  is  not 
so  striking  as  with  the  chloride. 

Praseodymium  chloride  in  mixtures  of  water  and  methyl  alcohol 
shows  the  same  general  features  as  were  manifested  by  the  chloride 
of  neodymium.  In  the  case  of  praseodymium  chloride  there  is  this 
additional  feature:  in  the  alcoholic  solution  an  entirely  new  band 
appears,  having  no  analogue  in  the  aqueous  solutions.  This  new  band 
in  the  ultra-violet  is  by  far  the  most  intense  in  the  entire  spectrum 
of  praseodymium  chloride.  On  adding  water  to  the  alcoholic  solution 
this  band  entirely  disappears.  In  this  case  the  alcohol  spectrum  is 
quite  different  from  the  water  spectrum. 

These  results  show  beyond  question  that  the  solvent  plays  an  impor- 
tant role  in  the  absorption  of  light  by  solutions.  The  question  arises, 
what  is  this  role?  It  is  difficult,  not  to  say  impossible,  to  explain  the 
action  of  the  solvent  on  any  other  ground  than  that  a  part  of  the  solvent 
combines  with  the  ions  and  molecules  of  the  dissolved  substance,  and 
the  solvated  parts  have  different  resonance  from  the  unsolvated.  This 
means  that  they  would  absorb  different  wave-lengths  of  light.  The 
alcoholates  would  have  different  resonance  from  the  hydrates,  whence 
the  different  absorption  spectrum  in  alcohol  from  that  in  water. 

We  regard  this  evidence  in  favor  of  solvation  in  solution  as  important, 
and,  as  we  shall  see,  many  examples  of  "solvent"  bands  were  brought 
to  light  in  the  investigation  which  followed. 

WORK  OF  JONES  AND  STRONG. 

The  work  of  Jones  and  Anderson  was  continued  by  Jones  and 
Strong.1  They  investigated  a  number  of  problems,  including  the  effect 
of  the  solvent  on  the  absorption  of  light  by  the  dissolved  substance. 
Jones  and  Anderson,  as  we  have  just  seen,  had  found  one  good  example 

Carnegie  Inst.  Wash.  Pubs.  Nos.  130  and  160.  Amer.  Chim.  Journ.,43,  37,  224  (1910);  45,  1 
(1910) ;  47,  27  (1912).  Phys.  Zeit.  10,  499  (1909).  Phil.  Mag.,  April,  1910.  Journ.  Chim.  Phys., 
8,  131  (1910). 


176  DISCUSSION   OF   EVIDENCE. 

of  the  existence  of  "solvent  bands"  in  the  absorption  spectra  of  neo- 
dymium  and  praseodymium  salts  in  water  and  the  alcohols.  The 
question  arose,  was  this  a  phenomenon  peculiar  to  these  salts,  or  does 
the  solvent  play  a  general  role  in  the  absorption  of  light  by  solutions? 

Jones  and  Strong  attempted  to  answer  this  question  by  studying  a 
large  number  of  salts  in  a  large  number  of  solvents.  They  worked 
especially  with  salts  of  neodymium  and  uranium,  because  these  sub- 
stances had  sharp  absorption  lines  and  bands  whose  positions  could 
easily  be  determined  with  reasonable  accuracy.  Work  was  done  not 
only  with  uranyl  salts,  but  with  uranous.  A  convenient  method  was 
found  for  reducing  uranyl  salts  to  the  uranous  condition,  and  uranous 
salts  were  found  to  have  very  sharp  absorption  lines. 

Uranyl  chloride  was  studied  in  the  following  solvents :  water,  methyl, 
ethyl,  propyl,  isopropyl,  butyl,  and  isobutyl  alcohols,  glycerol,  ether, 
methyl  ester,  and  formamide.  A  comparison  of  the  wave-lengths  of 
the  absorption  lines  and  bands  in  these  different  solvents  brought  out 
the  fact  that  the  wave-lengths  of  some  of  the  lines  and  bands  differed 
considerably  in  the  different  solvents.  The  results  here  showed  that 
the  solvent  unquestionably  has  much  to  do  with  the  absorbing  power 
of  the  solution,  "solvent  bands"  appearing  very  frequently.  The 
wave-lengths  of  a  few  of  the  different  lines  and  bands  of  uranyl  chloride 
in  the  above-named  solvents  have  been  tabulated,1  and  the  table  is 
here  reproduced.  It  shows  at  a  glance  the  different  wave-lengths  of 
the  several  lines  and  bands  compared. 

TABLE  51. — Wave-lengths  of  uranyl  chloride  absorption  lines. 

In  water XX  4025,      4170,     4315,       4460,     4560,     4740,  and  4920 

In  methyl  alcohol...     XX  4090,      4220,      4345,        4465,      4590,      4760,  and  4930 


In  ethyl  alcohol 

In  propyl  alcohol 

In  isopropyl  alcohol. . 

In  butyl  alcohol 

In  isobutyl  alcohol. 


XX  4100,      4250 4400,  4580,  4750,  and  4900 

XX  4100,      4230 4400,  4580,  4750,  and  4910 

XX  4100,      4250 4360,  4560,  4750 

XX  4100,      4240,      4390,  4560,  4750,  and  4970 

XX  4400,  4560,  4720,  and  4900 


In  ether XX  4040,      4160,      4300,  4444,      and  4630 

In  methyl  ester XX  4030,      4160,      4280,        4440,  4620,      4790,  and  4920 

In  glycerol XX  4025,      4140,      4260,        4400,  4540,      4720,  and  5050 

In  formamide XX  4450,  4650 and  4840 

The  absorption  spectra  of  uranyl  nitrate  in  mixtures  of  water  and 
methyl  alcohol  were  studied.  The  absorption  in  water  was  much  less 
than  in  pure  methyl  alcohol.  The  addition  of  water  to  the  alcoholic 
solution  diminished  the  absorption.  In  the  mixtures  of  water  and 
methyl  alcohol  the  absorption  bands  became  very  broad.  A  study 
of  these  broadened  bands  showed  that  they  were  the  "alcohol"  and 
"water"  bands  coexisting,  and  that  one  set  of  bands  was  not  simply 
the  other  set  shifted  in  position.  The  importance  of  this  fact  has 
already  been  referred  to  in  the  work  of  Jones  and  Anderson.  It  shows 
that  the  "alcohol"  bands  are  fundamentally  different  from  the  "water" 

Uourn.  Franklin  Inst.,  Dec.  1913,  p.  528;  also  Phil.  Mag.,  May  1912,  p.  730. 


DISCUSSION   OF  EVIDENCE.  177 

bands.  Further,  the  intensity  of  the  solvent  bands  is  a  function  of 
the  relative  amounts  of  the  solvents  that  are  present  in  the  mixture. 
This,  as  has  been  pointed  out,  indicates  the  existence  of  hydrates  in 
the  aqueous  solutions  and  of  alcoholates  in  solutions  in  alcohol,  these 
solvates  having  definite  resonance  and,  therefore,  definite  absorption 
spectra. 

One  of  the  most  striking  examples  of  solvent  bands  is  shown  by  the 
absorption  spectra  of  uranous  chloride  and  bromide  in  a  mixture  of 
water  and  methyl  alcohol.  We  find  two  entirely  distinct  spectra,  one 
belonging  to  each  solvent.  Some  lines  and  bands  appear  in  the  one  sol- 
vent which  are  entirely  absent  from  the  other,  and  practically  all  the 
lines  and  bands  have  very  different  positions  in  the  two  solvents.  To 
see  how  differently  the  spectra  appear,  reference  must  be  made  to  plate 
23  of  Publication  No.  160  of  the  Carnegie  Institution  of  Washington. 

The  spectrum  of  uranous  chloride  in  water  is  not  only  different  from 
the  spectrum  in  methyl  alcohol,  but  these  are  both  different  from  the 
spectrum  in  acetone.  If  we  compare  the  spectra  of  this  salt  in  the 
three  solvents,  we  might  easily  conclude  that  we  were  dealing  with 
three  fundamentally  different  spectra,  and  the  only  change  is  in  the 
nature  of  the  solvent. 

Uranous  salts  in  solvents  other  than  the  above  also  show  very 
characteristic  " solvent"  bands.  When  ethyl  alcohol  is  added  to  an 
aqueous  solution  of  uranous  chloride,  a  marked  change  is  produced 
in  the  spectrum.  The  "ethyl  alcohol"  bands  are  quite  different 
from  the  "water"  bands.  The  alcohol  bands,  or  the  water  bands, 
can  be  made  the  more  intense  by  simply  varying  the  relative  propor- 
tions of  the  two  solvents.  The  addition  of  acetone  to  an  aqueous 
or  methyl  alcohol  solution  of  uranous  chloride  produces  a  marked 
change  in  the  spectra.  A  number  of  acetone  bands  appear,  these 
being  different  from  the  "water"  bands  on  the  one  hand,  and  from 
the  "alcohol"  bands  on  the  other. 

Uranous  chloride  dissolved  in  methyl  alcohol  has  an  absorption 
spectrum  very  similar  to  that  in  ethyl  alcohol.  This  would  be  expected, 
on  account  of  the  close  similarity  of  methyl  alcohol  and  ethyl  alcohol. 
The  methyl  alcohol  bands  are  of  slightly  shorter  wave-lengths. 

The  absorption  spectra  of  uranous  chloride  in  glycerol,  and  in  mix- 
tures of  glycerol  and  water  were  also  studied.  A  number  of  "  glycerol " 
bands  manifested  themselves,  the  glycerol  absorption  being  very  dif- 
ferent from  that  of  water. 

The  absorption  spectrum  of  uranous  chloride  in  methyl  alcohol  and 
ether  was  also  studied.  The  solution  in  methyl  alcohol  showed  com- 
plete absorption  in  the  ultra-violet  to  wave-length  X  3700,  while  the 
addition  of  ether  extended  the  absorption  to  X  3800.  The  addition 
of  the  ether  caused  the  absorption  to  shift  towards  the  red,  the  magni- 
tude of  this  shift  being  from  10  to  30  A.  u. 


178  DISCUSSION   OP   EVIDENCE. 

It  has  already  been  pointed  out  that  salts  of  neodymium  are  espe- 
cially well  adapted  to  the  study  of  "solvent"  bands,  on  account  of  the 
sharpness  of  the  neodymium  lines  and  bands,  and  the  accuracy  with 
which  they  can  be  measured.  Neodymium  salts  were  studied  in  a 
number  of  solvents,  and  a  few  of  the  results  obtained  are  given  below.1 

ABSORPTION   SPECTRA   OF   NEODYMIUM   SALTS. 

The  following  nomenclature  will  be  used  in  describing  the  neodym- 
ium absorption  spectra: 

a  group  in  the  region  X3400  to  X3600. 

|8      "     at  about  X4300. 

7      "     from  X4600  to  X4800. 

5      "     from  X5000  to  X5400. 

€       "in  the  region  X5800. 

In  designating  the  neodymium  spectra  we  start  from  the  violet  end 
of  the  spectrum.  This  is  the  natural  method  when  a  grating  is  used. 
It  is  doubtful  whether,  in  the  near  future,  the  ultra-violet  spectrum 
of  neodymium  can  be  studied  much  farther  than  we  have  done,  so  that 
this  is  the  natural  end  of  the  spectrum  at  which  to  begin.  It  is,  on 
the  other  hand,  probable  that  there  are  many  neodymium  bands 
farther  down  in  the  infra-red  than  we  have  gone;  and  when  these  have 
been  worked  out  they  can  then  be  named  in  the  natural  order. 

The  change  in  the  absorption  spectrum  of  neodymium  chloride  as 
the  solvent  is  changed  can  best  be  seen  by  expressing  the  results  in 
the  following  form:  The  abbreviations  used  are  "d. "  diffuse,  "fa." 
faint,  "fi."  fine,  "h."  hazy,  "i."  intense,  "n."  narrow,  "sh."sharp, "st." 
strong,  "we."  weak,  "wi."  wide. 

The  following  results  obtained  with  neodymium  chloride  show  the 
effect  of  the  solvent  on  the  absorption  spectra  of  solutions  of  this 
compound.  The  bands  of  the  different  solvents  have  different  wave- 
lengths and  different  relative  intensities. 

Having  found  that  the  solvent  played  an  important  part  in  deter- 
mining the  absorption  of  light  by  the  dissolved  substances,  Jones  and 
Strong  used  isomeric  organic  solvents,  to  see  whether  such  closely  related 
compounds  would  affect  differently  the  power  of  substances  dissolved 
in  them  to  absorb  light.  They  prepared  solutions  of  neodymium 
chloride  in  propyl  and  isopropyl  alcohols,  and  in  butyl  and  isobutyl 
alcohols,  and  photographed  the  absorption  spectra  of  this  salt  in  these 
isomeric  solvents.  The  results  show  different  absorption  lines  and 
bands  in  the  isomeric  solvents. 

If  we  compare  carefully  the  spectra  of  neodymium  chloride  in  butyl 
and  isobutyl  alcohols,  we  find  that  the  bands  are  weak  and  diffuse 
in  isobutyl  alcohol,  and  have  different  relative  intensities  from  what 
they  have  in  the  butyl  alcohol.  The  bands  in  butyl  alcohol  are  very 

1See  Phil.  Mag.  May  1912,  p.  737,  from  which  the  few  following  pages  are  taken;  also,  Journ. 
Franklin  Inst.  Dec.  1913,  p.  531. 


DISCUSSION   OF  EVIDENCE. 


179 


TABLE  52.— Absorption  spectra  ofneodymium  chloride  in  certain  solvents, 
a  GROUP. 


In  water. 

In  methyl 
and  ethyl 
alcohols. 

In  propyl 
alcohol. 

In 
isopropyl 
alcohol. 

In  butyl 
alcohol. 

In  isobutyl 
alcohol. 

In 
glycerol. 

XX 
3390  we. 
3465  n.  st. 
3505  n.  st. 
3540  n.  st. 
3560 

XX 
3475  fa. 
3505 

3560  wi.i. 

XX 
3545  sh. 
3460 
3490 
3510  we. 
3525  st 

XX 
3460 
3510 
3535 

XX 
3450  sh.  n. 
3460  we. 
3492  d. 
3535  sh.  n. 
3545 

XX 
3455  we. 
3485  st. 
3515  we. 
3545 
3570 

XX 
3520  we. 
3475  st. 
3550  st. 

3540  st.  n. 

3560  d 

3560  we. 

3580  we. 

0  GROUP. 

4271  sh. 
4290  n.  we. 

4290 
4325 

4270  we. 

4285 

4265 

4265 

4285 

4300  we. 

4288  sh. 
4270  fi 

4330  wi  we 

4300 

4305  fi 

4450  wi.  we. 

-/  GROUP. 

4610  h. 

4700 

4600  we.  d. 

4600  d 

4700 

4620 

4645  we. 
4685 

4780 
4825 

4700 
4770 

4690 
4730 

4730 
4780 

4710 
4730 

4755  sh. 

4830 

4830 

4760 

4820  wi. 

4880  we. 

4790 

4840 

6  GROUP. 

5090  n. 
5125  wi.  h. 
5205  i.  n. 
5222  i  n 

5125  h. 
5180  h.  fa. 
5220  i.  n. 
5245  i 

5130  wi.  d. 
5180  wi.  d. 
5220 
5230 

5100  wi.  d. 
5320  wi.  d. 

5085  n. 
5095  n.  we. 
5130 
5200 

5150 
5260 
5215 
5230 

5120  wi.  h. 
5170  n. 
5190  n. 
5230 

5255  n 

5290  n 

5250 

5215 

5250 

5240 

5315  fa  h 

5315  fa 

5290 

5240 

5300 

5250 

5330  we. 

5270 
5300 

5270  we. 

e  GROUP. 

5725  n.  st. 
5745  st. 
5765  st, 
5795 

5725  h. 
5765  n. 
5800  st. 
58351. 
5860  h. 
5895  fa. 

5740 

5780 
5810 
5850 

5720  d. 
5780 
5810 

5750 

5780 
5820 
5860 
5900 
5930 

5740  we. 
5810  st, 
5850  st. 
5890 
5920 
5950  we. 
5995  we. 

5740  h. 
5790 
5805 
5820 
5850 

6920  we. 

180  DISCUSSION   OF   EVIDENCE. 

much  finer  and  sharper  than  they  are  in  isobutyl  alcohol.  Further, 
the  bands  of  neodymium  chloride  in  isobutyl  alcohol  have  slightly 
greater  wave-lengths  than  in  butyl  alcohol. 

To  eliminate  the  possibility  of  the  effect  of  the  solvent  on  absorption 
spectra  being  due  to  anything  inherent  in  the  nature  of  neodymium 
chloride,  the  nitrate  of  neodymium  was  studied  in  the  same  way  as 
the  chloride. 

The  absorption  spectra  of  neodymium  nitrate  in  water,  in  methyl 
alcohol,  in  ethyl  alcohol,  in  mixtures  of  these  alcohols  and  water,  in 
propyl  and  isopropyl  alcohols,  in  butyl  and  isobutyl  alcohols,  in  acetone 
and  in  mixtures  of  acetone  and  water,  in  ethyl  ester  and  in  formamide, 
were  carefully  photographed  and  studied.  Results  are  given  below  in 
the  case  of  neodymium  nitrate  only  for  the  a  bands. 

a  Bands. 

In  water. — Practically  the  same  as  the  bands  of  neodymium  chloride,  but 
the  bands  of  the  nitrate  are  broader  and  hazier  than  those  of  the  chloride. 

In  methyl  and  ethyl  alcohols. — There  are  only  two  bands  in  the  a  group, 
X  3465  and  X  3545. 

In  propyl  alcohol— -XX  3455,  3500,  and  3585. 

In  isopropyl  akohol.—\\  3460,  3505,  and  3535. 

In  butyl  alcohol— \\  3450,  3500,  and  3540. 

In  isobutyl  alcohol. — Ultraviolet  absorption  was  so  great  that  on  the  plate 
taken  the  a  group  did  not  appear.  The  absorption  in  general  is  the  same  as 
that  of  the  chloride  in  this  alcohol. 

In  acetone.— \\  3475,  and  3555. 

In  ethyl  ester— \\  3455,  3500,  and  3540. 

The  other  groups  of  absorption  bands  of  neodymium  nitrate  in  the 
different  solvents  show  differences  in  the  wave-lengths  comparable  with 
the  above;  but  these  results  suffice  to  show  the  effect  of  the  solvent 
on  the  power  of  neodymium  nitrate  to  absorb  light. 

The  above  is  strong  evidence  that  the  solvent  plays  an  important 
part  in  the  absorption  of  light  by  substances  dissolved  in  it.  When  we 
take  into  account  the  number  of  salts  studied  and  the  number  of 
solvents  employed,  the  evidence  is  little  short  of  proof.  The  only 
reasonable  question  is,  How  are  we  to  interpret  these  facts?  Before 
attempting  to  answer  this  question  we  should  take  into  account  also 
the  following  fact :  A  salt  dissolved  in  a  given  solvent  is  characterized 
by  a  definite  absorption  spectrum.  When  a  salt  is  dissolved  in  mix- 
tures of  varying  proportions  of  two  solvents,  only  two  definite  absorption 
spectra  appear,  one  being  characteristic  of  each  solvent.  One  spectrum 
does  not  gradually  change  into  the  other  as  the  composition  of  the  mixed 
solvent  changes,  but  only  the  relative  intensities  of  the  two  spectra 
vary.  Starting  with  that  mixture  of  the  two  solvents  in  which  both 
of  the  spectra  are  equally  intense,  if  we  diminish  the  amount  of  a 
relative  to  b,  the  spectrum  corresponding  to  a  becomes  feebler  and 
feebler,  and  the  spectrum  corresponding  to  b  more  and  more  intense. 


DISCUSSION   OP  EVIDENCE.  181 

This  fact  was  first  noted  by  Jones  and  Anderson,  and  since  repeatedly 
confirmed  by  the  work  of  Jones  and  Strong.  We  found  that  when 
neodymium  chloride  was  dissolved  in  a  mixture  of  methyl  alcohol  and 
water,  it  showed  a  definite  set  of  "water"  bands  and  a  definite  set  of 
"methyl  alcohol"  bands.  As  the  amount  of  water  in  the  solution 
was  decreased  relative  to  the  alcohol,  the  "water"  bands  decreased 
in  intensity  but  remained  in  the  same  position.  As  the  amount  of 
alcohol  in  the  solution  was  decreased  relative  to  the  water,  the  "alcohol " 
bands  decreased  in  intensity,  but  their  position  remained  unchanged. 

Jones  and  Anderson  interpreted  these  facts  as  strong  evidence  in 
favor  of  the  view  that  there  are  definite  hydrates  and  definite  alco- 
holates  in  the  solution. 

The  spectroscopic  evidence  for  solvation  in  solution  furnished  by 
Jones  and  Anderson  has,  as  has  already  been  stated,  been  increased 
many  fold  by  the  work  of  Jones  and  Strong.  A  large  number  of 
solvents  and  a  fairly  large  number  of  salts  have  been  used,  and  the 
existence  of  solvent  bands  in  general  has  been  established. 

The  question  of  the  relative  stability  of  the  different  solvates  with 
respect  to  various  physical  and  chemical  agents  has  been  studied  at 
length  by  Jones  and  Strong  by  means  of  absorption  lines  and  bands. 
It  would  lead  us  beyond  the  scope  of  this  paper  to  discuss  these  results 
in  detail.  Suffice  it  to  say  that  the  hydrates  in  general  are  the  most 
persistent  of  all  the  solvates,  although  this  depends  upon  the  conditions 
to  which  the  solution  is  subjected. 

Taking  all  of  the  spectroscopic  work  into  account,  I  regard  the  evi- 
dence from  this  source  as  strongly  supporting  the  solvate  theory  of 
solution  as  advanced  by  myself  about  fifteen  years  ago. 

EFFECT  OF  RISE  IN  TEMPERATURE. 

Jones  and  Strong  studied  the  effect  of  rise  in  temperature  on  the 
absorption  spectra  of  solutions.  Considerable  work  had  already  been 
done  on  the  effect  of  temperature  on  absorption  spectra  over  the  tem- 
perature range  0°  to  100°.  This  temperature  limit  could  be  studied 
in  open  vessels.  To  work  at  higher  temperatures  closed  apparatus 
must,  of  course,  be  used.  Such  apparatus  was  devised  and  used  up 
to  2000.1 

The  general  effect  of  rise  in  temperature  is  to  increase  the  color  of 
the  solution  of  the  inorganic  salt,  the  solution  becoming  less  trans- 
parent. The  deepening  of  the  color  is  usually  due  to  a  widening  of 
the  absorption  bands.  The  widening  of  the  bands  with  rise  in  tem- 
perature is  frequently  unsymmetrical. 

While  the  effect  of  rise  in  temperature  is  to  cause  the  long  wave- 
length bands  to  increase  in  intensity,  and  in  some  cases  to  produce 
new  bands,  in  some  solvents  the  effect  of  rise  in  temperature  is  to 
cause  the  short  wave-length  bands  to  increase  in  intensity  and  even 

Carnegie  Inst.  Wash.  Pub.  No.  130. 


182  DISCUSSION   OF   EVIDENCE. 

to  disappear.  If  the  absorption  is  sufficiently  intense  so  that  each 
group  of  bands  appears  as  a  single  band,  these  broad  bands  may  widen 
very  unsymmetrically  towards  the  red  as  the  temperature  is  raised. 

In  pure  solvents  the  bands  not  only  widen  with  rise  in  temperature, 
but  the  edges  become  more  diffuse.  With  mixtures  of  salts  such  as 
calcium  and  neodymium  chlorides,  the  bands  become  weaker  with  rise 
in  temperature. 

It  is  interesting  to  note  that  the  absorption  of  a  salt  in  mixtures  of 
two  solvents  often  decreases  in  intensity  with  rise  in  temperature. 
The  effect  of  rise  in  temperature  on  the  different  "solvent  bands"  is 
often  quite  different.  Uranous  bromide  in  40  per  cent  water  and  60 
per  cent  alcohol  showed,  at  ordinary  temperatures,  the  "water"  and 
1 '  alcohol "  bands  of  equal  intensity.  When  the  temperature  was  raised 
to  80°  the  "water"  bands  practically  disappeared,  while  the  "alcohol" 
bands  were  scarcely  widened  at  all. 

While  the  effect  of  rise  in  temperature  is  to  produce  a  change  in  the 
intensity  of  the  "solvate"  bands,  it  produces  very  little  change  in  the 
wave-lengths  of  the  bands.  In  some  cases,  in  mixtures  of  water  and 
alcohol,  the  alcohol  bands  increase  in  persistency  as  the  temperature 
is  raised.  This  is  important  as  showing  that  the  hydrates  and  alcohol- 
ates  have  different  degrees  of  stability  with  respect  to  temperature. 
The  effect  of  rise  in  temperature  is,  in  general,  to  increase  the  absorp- 
tion in  the  short  wave-lengths. 

An  interesting  question  arose  in  connection  with  the  effect  of  tem- 
perature on  the  solvates.  Does  rise  in  temperature  produce  a  perma- 
nent change  in  the  composition  of  the  solvates?  It  would  seem  highly 
probable  that  it  would  not.  The  composition  of  any  given  solvate  is 
determined  by  the  amount  of  solvent  relative  to  dissolved  substance. 
If  the  complex  solvate  is  rendered  simpler  by  rise  in  temperature,  then, 
when  the  solution  cools  down  the  solvate  should  have  its  original 
complexity — the  original  condition  of  equilibrium  should  be  restored. 

The  results  from  absorption  spectra  confirm  the  above  conclusion. 
When  a  salt  is  dissolved  in  a  mixture  of  two  solvents  and  the  solution 
heated,  there  is,  as  we  have  seen,  a  change  in  the  spectra.  When  the 
solution  is  cooled  again  the  original  spectrum  is  obtained.  This  shows 
that  the  original  solvates  are,  as  we  would  expect,  reformed.  This 
would  seem  to  have  some  bearing  on  the  nature  of  the  solvates  existing 
in  solution.  The  idea  that  was  originally  advanced  as  to  the  composi- 
tion of  the  hydrates  existing  in  aqueous  solution  was  that  a  large 
number  of  hydrates  existed  simultaneously  in  a  given  solution,  the 
composition  for  any  given  substance  and  any  given  solvent  being 
determined  chiefly  by  the  concentration  of  the  solution,  temperature 
being  constant.  In  a  word,  we  had  simply  a  condition  of  equilibrium 
for  any  given  substance  between  the  combined  and  the  free  water. 
This  condition  of  equilibrium  would  be  changed  with  rise  in  tempera- 
ture, some  of  the  combined  water  being  set  free,  in  accordance  with 


DISCUSSION   OF  EVIDENCE.  183 

the  general  principle  that  rise  in  temperature  breaks  down  aggregates 
formed  with  evolution  of  heat,  and  most  aggregates  are  formed  with 
heat  evolution. 

If  the  original  temperature  is  restored,  the  original  conditions  of 
equilibrium  are  reestablished  and  the  initial  solvates  reformed.  If 
there  were  only  a  few  definite  hydrates  in  any  given  aqueous  solution, 
each  of  these  would  probably  be  stable  over  a  definite  range  in  tem- 
perature, and  the  changes  in  their  composition  would  probably  take 
place  by  jumps.  This  would  produce  correspondingly  irregular  changes 
in  the  absorption  spectra,  and  not  the  regular  transitions  which  were 
noted. 

SPECTROPHOTOGRAPHY   OF   CHEMICAL   REACTIONS. 

The  effect  of  adding  an  acid  to  uranium  salts  of  another  acid  was 
studied  at  some  length.  Thus,  uranyl  nitrate  was  treated  with  sul- 
phuric, hydrochloric,  and  acetic  acids;  uranous  and  uranyl  acetates 
with  various  acids;  a  number  of  uranous  salts  and  neodymium  acetate 
with  nitric  acid,  and  so  on.  The  salts  and  acids  were  selected  so  as 
to  show  the  greatest  spectroscopic  changes.  The  action  of  nitric  acid 
on  uranous  salts  is  especially  interesting. 

The  spectrophotographs  of  chemical  reactions  show  that,  as  the 
salt  of  one  acid  is  transformed  into  the  salt  of  another  acid,  the  changes 
produced  in  the  spectra  are  gradual.  For  example,  when  uranyl 
nitrate  is  transformed  into  uranyl  sulphate,  the  uranyl  nitrate  bands 
gradually  shift  into  the  sulphate  position.  The  details  and  data 
bearing  on  this  point  are  given  in  Publication  No.  130  of  the  Carnegie 
Institution  of  Washington.  The  addition  of  a  large  amount  of  sul- 
phuric acid  to  a  small  amount  of  a  solution  of  uranyl  nitrate  in  nitric 
acid,  showed  admirably  the  gradual  shift  of  the  bands  from  the  nitrate 
to  the  sulphate  position. 

The  addition  of  a  small  amount  of  nitric  acid  to  uranous  acetate, 
does  not  appreciably  oxidize  the  uranous  salt.  The  uranous  bands 
are  shifted  towards  the  violet. 

The  gradual  shift  of  the  absorption  bands  as  one  salt  of  a  metal  is 
transformed  into  another  salt  by  the  addition  of  more  and  more  free 
acid  is  very  important. 

The  work  done  in  my  laboratory,  which,  up  to  the  time  we  are  now 
discussing  had  had  to  do  with  about  5,000  solutions,  had  shown  that 
any  given  series  of  absorption  bands  corresponds  to  a  definite  chemical 
condition  of  the  dissolved  substance.  When  a  salt  is  treated  with 
acid,  the  absorption  bands  of  some  of  the  salts  shift  gradually  over  to 
the  position  occupied  by  the  bands  corresponding  to  the  new  salt  of 
the  metal  with  the  acid  in  question.  In  such  a  case  the  absorption 
bands  can  be  made  to  occupy  any  position  between  the  initial  and 
final  positions.  It  therefore  seems  probable  that,  when  a  salt  of  one 
acid  is  transformed  in  this  way  into  a  salt  of  another  acid,  a  series  of 


184  DISCUSSION   OF   EVIDENCE. 

intermediate  systems  or  compounds  is  formed.  These  systems  are  for 
the  most  part  too  unstable  to  be  isolated,  at  least  by  the  methods  now 
at  our  disposal;  but  the  action  of  solutions  on  light  makes  their  exist- 
ence highly  probable. 

Our  chemical  equations  of  to-day  represent,  in  general,  only  the 
beginning  and  end  of  chemical  relations.  They  tell  us  little  or  nothing 
about  the  intermediate  stages  of  chemical  reactions,  and  these  are  the 
most  interesting  phases  of  the  reaction.  From  our  spectroscopic  work 
we  are  forced  to  conclude  that  at  least  some  chemical  reactions  are 
far  more  complex  than  would  be  indicated  by  the  equations  that  we 
ordinarily  use  to  express  them.  When,  for  example,  a  nitrate  is 
transformed  into  a  sulphate,  there  seems  to  be  formed  a  series  of  inter- 
mediate systems,  sulphonitrates  or  nitrosulphates.  We  know  nothing 
about  these  substances  chemically,  but  their  existence  is  made  highly 
probable  by  a  purely  physical  method — the  action  of  these  substances 
on  light. 

This  raises  the  question,  are  chemical  reactions  in  general  more  com- 
plex than  we  ordinarily  represent  them  to  be?  Do  these  intermediate 
systems  exist  in  chemical  reactions  in  general?  It  is  impossible  to 
study  all  reactions  by  the  spectroscopic  method,  if,  for  no  other  reason, 
because  many  solutions  do  not  have  sharp  and  well-defined  absorption. 
The  reactions,  however,  which  can  be  studied  directly  by  the  spectro- 
scopic method  do  not  seem  to  differ  in  any  fundamental  manner  from 
those  reactions  which  can  not  be  so  studied.  They  conform  to  the 
same  laws  that  are  obeyed  by  other  reactions  and  are  in  every  respect 
analogous  to  them.  This  leads  to  the  conjecture  that  in  those  reac- 
tions which  can  not  be  studied  spectroscopically,  there  are  also  inter- 
mediate systems  or  compounds  which  are  too  unstable  to  isolate;  and 
since  they  do  not  have  characteristic  spectra,  their  presence  can  not 
even  be  detected.  The  formation  of  these  intermediate  systems  is 
strictly  in  accord  with  the  action  of  mass  in  chemistry.  Some  such 
intermediate  compounds  have  in  a  number  of  reactions  recently  been 
isolated  by  methods  now  at  our  disposal.  As  methods  become  more 
refined,  and  we  acquire  better  control  of  conditions,  it  seems  not 
improbable  that  many  more  intermediate  compounds  will  be  isolated. 

At  present  we  can  not  isolate  any  large  percentage  of  these  inter- 
mediate systems  on  account  of  their  instability.  The  best  we  can  do 
is  to  study  their  properties  in  solution  in  the  different  solvents  by 
purely  physical  methods.  It  is  obvious  that  these  intermediate  sys- 
tems must  be  studied  if  we  are  ever  to  know  the  real  mechanism  of 
chemical  reactions,  and  not  simply  the  conditions  at  the  beginning  and 
end  of  reactions.  Other  suggestions  which  have  been  offered  to  explain 
the  gradual  shift  of  the  absorption  bands  as  one  salt  is  transformed  into 
another,  appear  to  be  entirely  inadequate,  if  not  meaningless. 


DISCUSSION   OP  EVIDENCE.  185 

WORK  OF  JONES  AND  GUY  ON  THE  ABSORPTION  SPECTRA  OF  SOLUTIONS. 

The  work  on  the  absorption  spectra  of  solutions  had,  at  the  time 
that  Guy  began  his  investigation,  been  extended  to  between  6,000  and 
7,000  solutions.  In  all  of  this  work  the  grating  spectroscope  had  been 
used,  and  the  results  recorded  on  a  photographic  plate.  The  photo- 
graphic method  recorded  the  positions  of  the  various  absorption  lines 
and  bands,  but  gave  only  a  qualitative,  or  at  best  a  roughly  quantita- 
tive indication  of  the  relative  intensities  of  the  various  lines  and  bands. 
The  photographic  method  is,  generally  speaking,  a  qualitative  method. 

If  we  are  ever  to  discover  relations  of  fundamental  significance 
between  the  power  of  dissolved  substances  to  absorb  light  and  the 
nature  of  solution,  we  must  have  some  quantitative  method  of  study- 
ing the  intensities  of  the  absorption  lines  and  bands  and  of  the  various 
parts  of  the  same  bands.  With  this  idea  in  mind  a  very  sensitive  radio- 
micrometer  was  built  and  used  to  measure  the  intensity  of  absorption. 

Before  taking  up  this  problem,  Jones  and  Guy  investigated  two  others 
by  the  photographic  method.  They  studied  the  effect  of  temperature 
on  the  absorption  spectra  of  aqueous  solutions  up  to  200°.  This 
required  a  specially  designed  apparatus  which  would  not  be  attacked 
by  the  superheated  water-vapor.  It  was  found  that  while  some  of  the 
bands  of  aqueous  solutions  are  practically  unaffected  by  rise  in  tem- 
perature, many  of  them  widen  as  the  temperature  is  raised.  The 
widening  of  the  absorption  bands  is  usually  not  symmetrical,  but  is 
generally  towards  the  red.  The  red  edge  widens  out,  becoming  more 
hazy  and  diffuse,  while  the  violet  edge  remains  pretty  sharp.  The 
effect  of  rise  in  temperature  on  the  absorption  spectra  of  aqueous  solu- 
tions is,  then,  often  analogous  to  the  effect  produced  by  increasing  the 
concentration  of  the  solution.  This  is  especially  the  case  with  solutions 
of  praseodymium  nitrate.  The  effect  of  dilution  on  absorption  spectra 
was  studied  pretty  thoroughly  by  Jones  and  Guy.  It  was  well  known 
that  both  molecules  and  ions  can  absorb  light,  and  the  question  was, 
do  they  have  the  same  or  different  absorption?  Jones  and  Anderson1 
had  shown  that  if  they  absorb  differently,  the  difference  is  slight.  To 
detect  any  such  differences  wide  ranges  in  dilution  must  be  employed. 

Cells  were  devised  for  holding  the  solutions,  which  were  0.5  cm., 
50  cm.,  and  250  cm.  in  length.  The  concentrations  were  varied  in 
the  same  proportions  as  the  lengths  of  the  cells.  If  we  call  the  con- 
centration used  in  the  shortest  cell  unity,  100  times  as  dilute  a  solution 
was  used  in  the  cell  which  was  50  cm.  long,  and  this  was  diluted  5  tunes 
for  the  longest  cell.  It  was  found  that  many  of  the  absorption  bands 
of  neodymium  chloride  and  bromide  widen  as  the  concentration  of  the 
solution  is  increased.  Some  of  the  bands  of  neodymium  sulphate  and 
acetate  show  similar  changes  with  increase  in  the  concentration  of  the 
solution.  The  most  marked  changes,  however,  are  produced  with  the 
bands  of  neodymium  nitrate.  Many  of  them  show  very  pronounced 

'Carnegie  Inst.  Wash.  Pub.  No.  110. 


186  DISCUSSION   OF   EVIDENCE. 

widening  with  increase  in  concentration.  Solutions  of  praseodymium 
salts  also  show  a  widening  of  the  absorption  bands  as  the  concentra- 
tions are  increased,  but  these  changes  are  less  pronounced  than  with 
salts  of  neodymium.  The  absorption  spectra  of  uranyl  salts  change 
more  with  change  in  concentration  than  the  spectra  even  of  salts  of 
neodymium.  The  changes  are  in  the  same  direction,  the  bands  increas- 
ing in  breadth  with  increase  in  concentration. 

These  results  are  what  would  be  expected  from  the  solvate  theory  of 
solution.  As  the  concentration  of  the  solution  is  changed,  the  com- 
plexity of  the  solvate  about  the  molecules  or  ions  is  changed.  It  would 
seem  that  this  ought  to  affect  the  resonance  of  the  solvated  resonator. 
As  the  concentration  of  the  solution  is  increased  the  solvate  becomes 
simpler  and  simpler.  The  vibrating  particle  surrounded  by  a  simple 
solvate  should  show  different  absorption  than  when  surrounded  by  a 
complex  solvate.  The  above  results  show  that  such  is  the  case,  the  more 
concentrated  the  solution  the  wider  in  general  the  absorption  bands. 

The  radiomicrometer  not  only  provides  us  with  a  method  of  study- 
ing absorption  spectra  quantitatively,  but  greatly  extends  the  range 
of  wave-lengths  that  can  be  studied.  The  earlier  work  with  the  very 
sensitive  radiomicrometer  had  to  do  with  the  study  of  solutions  of 
neodymium  salts.  The  effect  of  dilution  on  absorption  spectra  was 
also  investigated  quantitatively  by  means  of  the  radiomicrometer. 
It  was  found  by  this  method,  as  with  the  grating  and  photographic 
plate,  that  the  more  concentrated  the  solution  the  broader  the  absorp- 
tion bands.  It  was  also  found  that  in  the  more  dilute  solution,  while 
the  absorption  bands  were  narrower,  they  were  more  intense.  Further, 
in  the  more  dilute  solutions  the  centers  of  the  bands  were  displaced 
towards  the  longer  wave-lengths. 

The  most  interesting  and  important  result  brought  out  by  the  work 
of  Jones  and  Guy  was  the  effect  of  the  dissolved  substance  on  the  ab- 
sorption of  light  by  water.  We  noted  that  aqueous  solutions  of  certain 
hydrated  salts  are  more  transparent  than  pure  water.  This  is  obvi- 
ously a  fact  which  called  for  careful  study.  We  compared  the  absorp- 
tion of  aqueous  solutions  of  strongly  hydrated  salts, with  the  absorption 
of  a  layer  of  water  equal  in  depth  to  the  water  in  the  solution.  Similar 
experiments  were  carried  out  with  salts  which  are  only  slightly 
hydrated.  The  slightly  hydrated  salts  with  which  we  worked  were 
potassium  chloride  and  ammonium  chloride  and  nitrate.  It  was 
necessary  to  select  colorless  salts  which  themselves  had  little  or  no 
absorption  in  the  infra-red  where  water  absorbs.  It  was  found,  in  the 
earlier  work,  that  the  above-named  compounds  had  nearly  the  same 
absorption  as  water  having  the  same  depth  as  the  water  in  the  solution; 
but  in  subsequent  work  this  conclusion  must  be  modified  for  certain 
substances  near  the  bottoms  of  the  absorption  bands. 

In  terms  of  the  solvate  theory  of  solution,  we  should  expect  the 
absorption  of  the  solution  of  a  slightly  hydrated  salt  in  general  not  to 


DISCUSSION   OF   EVIDENCE.  187 

differ  very  greatly  from  that  of  so  much  pure  water,  since,  when  the 
solvent  is  not  combined  with  the  dissolved  substance,  it  is  difficult  to  see 
how  either  could  affect  appreciably  the  absorbing  power  of  the  other. 

When  we  turned  to  the  strongly  hydrated  salts,  very  different  results 
were  obtained.  As  examples  of  this  class  of  substances  we  studied 
calcium  and  magnesium  chlorides  and  aluminium  sulphate.  Take  the 
results  for  a  5.3  normal  solution  of  calcium  chloride.  The  solution  is 
more  transparent  from  0.9  n  to  1  p..  It  is  again  the  more  transparent 
from  1.05ju  to  1.2/i,  being  as  much  as  25  per  cent  more  transparent 
than  the  solution.  For  the  longer  wave-lengths  the  water  is  in  general 
the  more  transparent  until  1.42M  is  reached,  when  both  water  and 
solution  become  equally  opaque.  Similar  results  were  obtained  with 
magnesium  chloride. 

Aluminium  sulphate  presents  this  peculiarity,  that  at  1  //  the  solu- 
tion is  more  transparent  than  the  water.  The  obvious  explanation  of 
these  surprising  results  seems  to  be  that  they  must  be  due  to  some 
action  of  the  dissolved  substance  on  the  solvent.  Jones  and  Anderson1 
showed  that  the  solvent  can  have  a  marked  effect  on  the  absorbing 
power  of  the  solution  in  that  solvent,  even  when  the  solvent  itself  had 
no  absorption  in  the  region  in  question. 

A  large  number  of  examples  of  "solvent  bands"  were  discovered  by 
Jones  and  Strong.2  They  found  many  non-absorbing  solvents  which 
affected  the  absorption  of  the  dissolved  substance,  and  could  even  dis- 
tinguish between  certain  organic  solvents  and  their  "iso"  compounds 
by  the  "solvent  bands"  which  manifested  themselves.  This  action 
seems  to  have  been  satisfactorily  explained  as  due  to  a  combination  of 
the  solvent  with  the  dissolved  substance  forming  solvates.  The  sol- 
vate  theory  enables  us  to  account  for  many  facts  which  apparently 
could  not  be  satisfactorily  explained  by  the  theory  of  electrolytic  dis- 
sociation alone,  as  we  have  seen.  The  same  theory  seems  to  aid  us  in 
explaining  the  facts  just  described.  Those  compounds  which  do  not 
form  hydrates,  or  which  form  only  very  simple  hydrates,  such  as 
potassium  chloride  and  the  like,  show  results  such  as  would  be  expected. 
Their  solutions  are  not  more  transparent  than  so  much  pure  water. 
In  general,  the  absorption  of  such  solutions  is  of  the  same  order  of 
magnitude  as  that  of  the  water  in  which  they  are  dissolved.  We  shall 
see  that  it  came  out  in  later  work  that  solutions  of  only  slightly  hy- 
drated salts  are  more  opaque  than  pure  water  at  the  centers  of  the 
absorption  bands.  This,  however,  does  not  affect  at  all  the  conclu- 
sions drawn  above.  It  is  only  the  hydrated  salts  whose  solutions  are 
appreciably  more  transparent  than  so  much  pure  water.  How  does 
the  solvate  theory  explain  these  facts? 

The  combined  water  seems  to  have  less  power  to  absorb  light  than 
free  water.  This  would  account  for  the  above  facts.  The  presence 

Carnegie  lust.  Wash.  Pub.  No.  110.  *Ibid.,  Nos.  130  and  160. 


188  DISCUSSION    OF   EVIDENCE. 

of  the  salt  seems  to  shift  the  absorption  of  the  water  towards  the 
larger  wave-lengths.  Rise  in  temperature  and  increase  in  concentra- 
tion shift  the  absorption  of  the  salt  towards  the  longer  wave-lengths. 
The  effect  of  rise  in  temperature  and  increase  in  concentration  is  to 
simplify  the  hydrates  existing  in  the  solution.  Simplifying  the  resona- 
tor, then,  shifts  the  absorption  towards  the  red. 

The  effect  of  the  salt  on  the  absorption  of  the  water,  is  the  same  as 
rise  in  temperature  and  increase  in  the  concentration  of  the  solution 
on  the  absorption  of  the  dissolved  substance.  It  may  well  be  that  the 
dissolved  substance  diminishes  the  association  of  the  solvent  and  this 
simplifies  the  solvent  resonator.  This  may  be  true,  especially  with 
water  of  hydration,  which  is  more  directly  under  the  influence  of  the 
dissolved  substance  than  the  free  water. 

WORK  OF  JONES.  SHAEFFER,  AND  PAULUS. 

The  result  obtained  by  Jones  and  Guy  was  regarded  as  of  such 
importance  in  its  bearing  on  the  solvate  theory  of  solution,  that  it 
was  thought  desirable  to  repeat  and  elaborate  with  improved  method 
the  work  which  led  to  it.  Certain  details  of  method  and  manipulation 
were  carefully  studied,  and  the  degree  of  accuracy  of  the  procedure 
adopted  was  carefully  ascertained.  This  has  all  been  discussed  in 
detail  in  the  first  chapter  of  this  monograph.  The  non-hydrating  or 
slightly  hydrating  salts,  potassium  chloride,  ammonium  bromide,  and 
sodium  nitrate,  were  studied.  The  strongly  hydrated  calcium  chloride, 
magnesium  chloride,  magnesium  bromide,  magnesium  sulphate,  mag- 
nesium nitrate,  zinc  sulphate,  and  zinc  nitrate  were  investigated  at 
varying  concentrations  and  depths  of  layers. 

Solutions  of  the  strongly  hydrated  salts  have  in  general  greater 
transparency  than  pure  water,  especially  at  the  centers  of  the  absorp- 
tion bands.  As  the  regions  of  intense  absorption  are  approached  in  the 
longer  wave-lengths,  the  solution  is  much  more  transparent  than  the 
pure  solvent.  This  difference  may  amount  to  as  much  as  40  per  cent. 

The  non-hydrated  or  only  slightly  hydrated  salts  give  results  which, 
in  many  respects,  are  exactly  the  opposite  of  those  obtained  with 
hydrated  salts.  In  the  three  cases  studied,  the  solution  had  greater 
absorption  than  the  solvent  at  the  centers  of  the  bands.  This  is  pre- 
cisely the  opposite  of  what  was  found  for  the  strongly  hydrated  salts. 
Regions  of  the  spectrum,  for  which  solutions  of  hydrated  salts  were  as 
much  as  40  per  cent  more  transparent  than  the  solvent,  show  for  non- 
hydrated  salts  that  the  solution  is  40  per  cent  less  transparent. 

It  was  pointed  out  that  the  results  obtained  could  be  best  explained 
by  the  solvate  theory  of  solution.  Indeed,  this  evidence  is  of  the  very 
strongest  for  that  theory.  In  the  solutions  studied,  more  than  half 
of  the  water  was  shown  to  be  combined  with  the  dissolved  substance. 
It  was  shown  that  this  would  certainly  alter  the  vibrational  frequency 
or  resonance  of  the  absorbing  systems. 


DISCUSSION   OF  EVIDENCE.  189 

The  transmission  curves  obtained  seem  to  justify  the  conclusion  that 
combined  water  has  less  power  to  absorb  light  than  uncombined.  We 
have  been  able  to  find  no  other  rational  explanation  which  would 
account  satisfactorily  for  our  results.  The  difference  in  the  behavior 
of  hydrated  and  non-hydrated  salts  seems  unquestionable. 

Any  attempt  to  explain  such  a  difference  as  the  above  on  the  ground 
of  a  change  in  the  dielectric  constant  of  the  medium  does  not  appear  to 
have  a  good  physical  basis.  Why  the  presence  of  the  one  class  of 
salts  alters  the  dielectric  constant  of  the  medium  differently  from  the 
other  class,  is  a  question  that  would  have  to  be  answered.  This 
attempt  to  explain  our  results  does  not  appear  to  be  much  more  than 
words.  We  regard,  then,  the  spectroscopic  evidence  in  its  bearing  on 
the  solvate  theory  of  solution  as  among  the  most  important.  The 
presence  of  definite  "solvent  bands"  in  the  different  solvents  and  the 
difference  between  the  absorption  of  aqueous  solutions  of  non-hydrated 
and  strongly  hydrated  salts  are  to  be  counted  as  among  the  strongest 
and  most  direct  lines  of  evidence  thus  far  brought  to  light  in  my  labora- 
tory bearing  on  the  solvate  theory  of  solution. 

SUMMARY  OF  THE  LINES  OF  EVIDENCE  OBTAINED  IN  THIS  LABORATORY  BEARING 
ON  THE  SOLVATE  THEORY  OF  SOLUTION. 

The  following  lines  of  evidence  bearing  on  the  solvate  theory  of 
solution  have,  then,  been  established  in  this  laboratory.1 

1.  Relation  between  lowering  of  the  freezing-point  of  water  and 
water  of  crystallization  of  the  dissolved  substance. 

2.  Approximate  composition  of  the  hydrates  formed  by  various 
substances  in  solution. 

3.  Relation  between  the  minima  in  the  freezing-point  curves  and 
the  minima  in  the  boiling-point  curves. 

4.  Relation  between  water  of  crystallization  and  temperature  of 
crystallization. 

5.  Hydrate  theory  in  aqueous  solutions  becomes  the  solvate  theory 
in  solutions  in  general. 

6.  Temperature  coefficients  of  conductivity  and  hydration. 

7.  Relation  between  hydration  of  the  ions  and  then*  ionic  volumes. 

8.  Hydration  of  the  ions  and  the  velocities  with  which  they  move. 

9.  Dissociation  as  measured  by  the  freezing-point  method  and  by 
the  conductivity  method. 

10.  Effect  of  one  salt  with  hydrating  power  on  the  hydrates  formed 
by  a  second  salt  in  the  same  solution. 

11.  Investigations  in  mixed  solvents. 

12.  Spectroscopic  evidence  bearing  on  the  solvate  theory  of  solution; 
work  of  Jones  and  Uhler.  m 

13.  Work  of  Jones  and  Anderson  on  absorption  spectra,  in  which 
the  presence  of  " solvate"  bands  was  first  detected.     This  showed  that 

'See  Journ.  Franklin  Inst.,  Dec.  1913. 


190  DISCUSSION   OF   EVIDENCE. 

the  solvate  had  an  effect  on  the  absorption  of  light,  and  this  could  be 
explained  only  as  due  to  a  combination  between  the  solvent  and  the 
resonator,  or  something  containing  the  resonator. 

14.  The  work  of  Jones  and  Strong  on  absorption  spectra  established 
the  existence  of  a  larger  number  of ' '  solvent  "bands .    They  showed  that 
these  were  formed  by  many  salts  and  in  many  solvents.     They  could 
even  distinguish  between  the  bands  of  a  salt  in  a  given  alcohol  and  in 
its  isomer.     This  was  regarded  as  very  important.     The  temperature 
work  of  Jones  and  Strong  was  strong  evidence  for  the  solvate  theory. 

15.  The  work  of  Jones  and  Guy  on  the  effect  of  high  temperature  on 
the  absorption  spectra  of  aqueous  solutions,  and  also  on  the  effect  of  dilu- 
tion, led  to  results  which  were  all  in  keeping  with  the  solvate  theory. 

The  most  important  spectroscopic  work  of  Jones  and  Guy,  which 
bears  on  the  solvate  theory  of  solution,  is  that  in  which  the  radio- 
micrometer  was  used.  It  was  here  shown  that  solutions  of  certain 
strongly-hydrated  non-absorbing  salts  are  more  transparent  than  pure 
water  having  a  depth  equal  to  that  of  the  water  in  the  solution.  In  the 
case  of  non-hydrated  salts  the  solution  was  the  more  opaque.  This 
shows  that  water  in  combination  with  the  dissolved  substance — water 
of  hydration — has  less  absorption  than  pure,  free  water.  This  is 
regarded  as  striking  evidence  that  some  of  the  water  in  the  presence 
of  salts  which  are  shown  by  other  methods  to  hydrate  is  different 
from  pure,  free,  uncombined  water;  and  the  simplest  explanation  seems 
to  be  that  this  is  the  combined  water,  or  the  water  of  hydration. 

16.  The  work  of  Jones  and  Guy  was  repeated  and  extended  by 
Jones,  Shaeffer,  and  Paulus.     They  obtained  results  of  the  same  general 
character  as  those  found  by  Jones  and  Guy.     Solutions  of  hydrated 
salts  were  in  general  more  transparent  than  pure  water,  especially  at 
the  centers  of  the  absorption  bands.     Solutions  of  non-hydrated  or 
only  slightly  hydrated  salts  are  more  opaque  than  pure  water,  especially 
at  the  centers  of  the  bands. 

The  above  sixteen  lines  of  evidence  all  point  to  the  general  correct- 
ness of  the  view  that  when  a  salt  is  dissolved  in  a  solvent  there  is  more 
or  less  combination  between  the  salt,  or  the  ions  resulting  from  it, 
and  the  solvent.  The  magnitude  of  this  solvation  depends  upon  the 
nature  of  the  substance  and  of  the  solvent. 

HOW  THE  PRESENT  SOLVATE  THEORY  OF  SOLUTION  DIFFERS  FROM  THE  OLDER 
HYDRATE  THEORY. 

The  present  solvate  theory  of  solution  is  not  simply  one  of  several 
possible  suggestions  which  accounts  for  a  certain  class  of  experimental 
facts.  It  is  the  only  suggestion  that  has  thus  far  been  made  which 
seems  to  account  satisfactorily  for  all  of  the  facts  established.  Most 
of  the  above  sixteen  lines  of  evidence  bearing  on  solvation  in  solution 
were  obtained  as  the  direct  result  of  experimental  work  suggested  by 
the  solvate  theory  and  carried  out  to  test  this  theory.  Many  of  the 


DISCUSSION   OF  EVIDENCE.  191 

results  were  predicted  from  this  theory  before  a  single  experiment 
was  carried  out.  Solvation,  then,  being  accepted,  as  now  seems  pretty 
generally  the  case,  the  question  arises,  how  does  the  present  solvate 
theory  of  solution  differ  from  the  older  hydrate  theory  of  Mendele"eff, 
which  has  long  since  been  abandoned  as  untenable? 

Mendele"erTs  theory  was  that  certain  hydroscopic  substances,  such 
as  calcium  chloride,  sulphuric  acid,  and  the  like,  formed  a  few  definite 
hydrates  when  in  the  presence  of  water.  Thus,  sulphuric  acid  formed 
the  hydrates  H2S04.2H2O,  H2S04.6H2O,  H2SO4.100H20. 

This  view  of  Mendele*eff  was  proposed  as  the  result  especially  of 
measuring  the  specific  gravities  of  aqueous  solutions  of  such  com- 
pounds at  different  dilutions.  When  the  specific  gravities  were  plotted 
against  the  concentrations,  the  curve  was  not  a  continuous  one,  but 
showed  a  number  of  breaks.  These  breaks  Mendele*eff  could  account 
for  by  assuming  that  certain  definite  hydrates  or  compounds  between 
water  and  the  dissolved  substances  existed  at  these  concentrations. 
This  was  among  the  most  important  evidence  brought  to  light  bearing 
on  the  so-called  hydrate  theory  of  Mendele"eff. 

This  suggestion  of  Mendele"eff,  based  upon  such  inadequate  evidence, 
should  not  be  called  a  theory.  It  is  scarcely  worthy  of  the  name 
hypothesis.  Before  a  suggestion  becomes  a  theory  there  should  be  a 
fair  amount  of  evidence  supporting  it,  and  showing  not  only  that  the  sug- 
gestion accounts  for  the  facts,  but  that  it  is  the  only  suggestion  which 
will  account  for  them.  This  was  lacking  in  the  so-called  Mendele"eff 
hydrate  theory. 

The  present  solvate  theory  of  solution  may  claim  to  have  a  fairly 
good  experimental  support,  as  the  above  review  of  the  evidence  ob- 
tained in  this  laboratory  will  show.  In  aqueous  solutions  hydration 
is  a  general  phenomenon.  Some  substances  combine  with  very  little 
water,  but  most  salts  combine  with  very  large  amounts  of  water,  the 
amount  of  combined  water  for  any  given  substance  being  a  function 
of  the  concentration  of  the  solution  and  of  the  temperature.  The 
more  dilute  the  solution  the  larger  the  amount  of  the  solvent  com- 
bined with  the  dissolved  substance — the  more  complex  the  hydrate. 
The  lower  the  temperature  the  more  complex  the  solvate.  These 
solvates  are  very  unstable;  indeed,  so  unstable  that  it  seems  better  to 
call  them  systems  than  definite  chemical  compounds.  Anything  so 
easily  broken  down  by  rise  in  temperature  could  hardly  be  called  a 
chemical  compound.  Here,  again,  the  present  solvate  theory  differs 
from  the  older  hydrate  theory. 

While  there  is  some  spectroscopic  evidence  pointing  to  the  existence 
in  solution  of  a  certain  definite  hydrate,  or  certain  definite  hydrates, 
we  have  obtained  a  large  amount  of  evidence  which  seems  to  indicate 
the  existence  in  aqueous  solutions  of  a  large  number  of  hydrates,  or 
indeed  of  a  whole  series  of  hydrates,  the  composition  depending  pri- 
marily on  the  concentration  of  the  solution.  While  this  is  not  essential 


192  DISCUSSION   OF   EVIDENCE. 

to  the  present  solvate  theory  of  solution,  it  would  differentiate  it  funda- 
mentally from  the  older  hydrate  theory. 

The  present  theory  is  not  simply  a  hydrate  theory  of  aqueous  solu- 
tions. Evidence  has  been  obtained,  and  is  herein  briefly  discussed, 
which  shows  that  solvents  other  than  water  combine  with  the  dis- 
solved substance.  This  has  been  established  for  the  alcohols  by  the 
boiling-point  method,  and  for  the  alcohols  and  many  other  solvents  by 
spectroscopic  investigations.  Indeed,  enough  evidence  has  already  been 
obtained  to  make  it  highly  probable  that  solvation  is  not  limited  to 
aqueous  solutions,  but  is  a  general  property  of  solutions.  Solvents  in 
general  have  more  or  less  power  to  combine  with  substances  dissolved  in 
them — in  a  word,  we  have  the  solvate  instead  of  simply  a  hydrate  theory. 

A  method  has  been  worked  out  in  this  laboratory  for  determining 
the  approximate  composition  of  the  hydrates  existing  in  aqueous  solu- 
tions. This  makes  the  present  theory  useful  scientifically.  We  can 
now  determine  approximately  the  amount  of  " combined"  and  the 
amount  of  "  free  "  water  existing  in  any  given  aqueous  solution.  Thus, 
our  theory  is  placed  upon  a  workable  basis,  and  enables  us  to  determine, 
in  any  given  case,  how  much  of  the  liquid  present  is  really  playing  the 
role  of  solvent. 

The  evidence  pointing  to  the  general  correctness  of  the  solvate  theory 
of  solution  is,  then,  so  strong  that  it  seems  that  this  conception  is  in 
accord  with  a  fundamental  condition  in  connection  with  the  nature  of 
solution. 

Further,  our  solvate  theory  of  solution  is  very  different  from  the 
earlier,  unproved  hydrate  theory  of  Mendele"eff. 

The  question  now  arises,  of  what  scientific  significance  or  value  is  the 
establishing  of  the  fact  that  there  is  more  or  less  combination  between 
the  dissolved  substance  and  the  solvent? 

SIGNIFICANCE  OF  THE  SOLVATE  THEORY  OF  SOLUTION.1 

The  evidence  for  the  solvate  theory  of  solution,  which  has  been  fur- 
nished in  this  laboratory  as  the  result  of  somewhat  more  than  a  dozen 
years  of  investigation,  has  recently  been  brought  together  and  briefly 
discussed.2  The  evidence  is  so  unambiguous  and  convincing,  that  ions 
and  some  molecules  combine  with  more  or  less  of  the  solvent,  that  it 
seems  that  it  can  now  be  accepted  as  a  fact  of  science. 

This,  however,  raises  a  number  of  questions :  What  relation  does  the 
solvate  theory  of  solution  bear  to  the  theory  of  electrolytic  dissociation? 

Does  the  solvate  theory  help  us  to  explain  any  of  the  apparent  dis- 
crepancies in  the  theory  of  electrolytic  dissociation?  Does  the  solvate 
theory  help  us  to  explain  the  facts  of  chemistry  in  general  and  of 
physical  chemistry  in  particular?  Why  is  the  nature  of  solution  so 
important,  not  only  for  chemistry  but  for  science  in  general? 

1This  section  is  taken  directly  from  my  paper  in  the  Journal  of  the  Franklin  Institute,  Dec.  1913. 
2Zeit.  phys.  Chem.,  74,  325  (1910). 


DISCUSSION   OF   EVIDENCE.  193 

THE  SOLVATE  THEORY  AND  THE  THEORY  OF  ELECTROLYTIC  DISSOCIATION. 

When  Arrhenius  proposed  the  theory  of  electrolytic  dissociation, 
the  question  was  not  even  raised  as  to  the  condition  of  the  ions  in  the 
solution,  except  that  they  behave  as  if  they  existed  independently  of 
one  another  in  solution.  The  theory  simply  said  that  molecules  of 
acids,  bases,  and  salts,  in  the  presence  of  a  dissociating  solvent  like 
water,  break  down  to  a  greater  or  less  extent  into  charged  parts  called 
ions,  the  cations  or  positively  charged  parts  being  electrically  equiva- 
lent to  the  anions  or  negatively  charged  parts.  The  cations  are  usually 
simple  metallic  atoms  carrying  one  or  more  unit  charges  of  positive 
electricity.  The  cation  might,  however,  be  more  or  less  complex,  as 
illustrated  by  ammonium  and  its  substitution  products.  The  anion  is 
usually  complex,  consisting  of  a  larger  or  smaller  number  of  atoms. 
It  may,  however,  be  an  atom  carrying  negative  electricity,  as  in  the 
case  of  the  halogen  acids  and  their  salts. 

The  degree  of  dissociation  is  determined  by  the  nature  of  the  acid, 
base,  or  salt.  Strong  acids  and  bases  are  greatly  dissociated.  Indeed, 
the  degree  of  dissociation  determines  their  strength.  Nearly  all  of  the 
salts  are  strongly  dissociated  compounds,  there  being,  however,  some 
exceptions,  as,  notably,  the  halogen  salts  of  mercury,  cadmium,  and 
zinc.  There  are,  however,  considerable  differences  in  the  amounts  to 
which  salts  in  general  are  dissociated  at  the  same  dilution. 

The  quantitative  evidence  furnished  by  Arrhenius  and  others  for  the 
theory  of  electrolytic  dissociation  is  so  convincing  that  few  chemists 
of  any  prominence,  who  have  carefully  examined  the  evidence,  have 
ever  doubted  the  general  validity  of  the  theory;  and  the  theory  has 
become  substantiated  by  such  an  abundance  of  subsequently  discovered 
facts  that  it  has  now  become  a  law  of  nature  and  a  fundamental  law 
of  chemical  science. 

Arrhenius  saw  and  pointed  out  clearly  the  importance  of  ions  for 
chemistry ;  Ostwald  and  his  pupils  have  shown  that  chemistry  is  essen- 
tially a  science  of  the  ion,  molecules  for  the  most  part  being  incapable 
of  reacting  chemically  with  molecules;  and  Nernst  has  proved  that  the 
ion  is  the  active  agent  in  all  forms  of  primary  cells. 

The  theory  of  electrolytic  dissociation,  as  already  stated,  does  not 
raise  the  question  as  to  the  relation  between  the  ion  and  the  solvent. 
At  the  time  that  the  theory  was  proposed,  chemists  did  not  know,  and 
probably  had  no  means  of  finding  out,  whether  the  ion  is  or  is  not  com- 
bined with  the  solvent  in  contact  with  it.  The  solution  of  this  problem 
remained  for  subsequent  work. 

Some  of  the  many  lines  of  evidence  that  ions  and  certain  molecules 
are  combined  with  a  larger  or  smaller  number  of  molecules  of  the  sol- 
vent, and  in  many  cases  with  a  very  large  number  of  molecules  of  the 
solvent,  have  been  recently  discussed  briefly  by  Jones  in  an  article  in 


194  DISCUSSION   OF  EVIDENCE. 

the  Zeitschrift  fur  physikalische  Chemie.1  The  amount  of  the  solvent 
combined  with  an  ion  is  primarily  a  function  of  the  nature  of  the  ion 
or  ions  in  the  solution.  It  is,  however,  conditioned  very  largely  by  the 
dilution  of  the  solution,  and  also  by  the  temperature. 

The  evidence,  some  of  which  is  given  in  the  paper  referred  to  above, 
and  the  remainder  in  other  publications,  of  the  results  of  investigations 
carried  out  in  this  laboratory  during  the  past  fifteen  years,  shows 
that  the  power  of  the  ions  to  combine  with  the  solvent  is  by  no  means 
limited  to  water  and  aqueous  solutions,  but  is  a  property  of  solutions 
in  general.  The  alcohols,  acetone,  glycerol,2  etc.,  combine  with  certain 
substances  dissolved  in  them;  and  it  seems  more  than  probable  that 
all  solvents  combine  with  the  dissolved  substances  to  a  greater  or  less 
extent.  In  a  word,  we  do  not  have  simply  a  theory  of  hydration,  but  a 
theory  of  solvation  in  general,  which  is  an  essential  part  of  any  general- 
ization that  would  take  into  account  the  facts  presented  by  solution. 

The  solvate  theory  of  solution  has  been  regarded  in  some  quarters 
as  a  rival  of  the  electrolytic  dissociation  theory  of  solution,  if  not 
directly  antagonistic  to  it.  Such  is  not  at  all  the  case.  The  solvate 
theory  begins  where  the  theory  of  electrolytic  dissociation  ends.  The 
latter  gives  us  the  ions  from  molecules,  and  the  former  tells  us  the  con- 
dition of  the  ions  in  the  presence  of  a  solvent  after  they  are  formed. 

The  solvate  theory  of  solution,  then,  simply  supplements  the  theory 
of  electrolytic  dissociation,  and  both  must  be  taken  into  account  if  we 
ever  wish  to  understand  the  phenomena  presented  by  solution. 

DOES  THE   SOLVATE   THEORY  HELP   TO   EXPLAIN   ANY  OF  THE  APPARENT 
EXCEPTIONS  TO  THE  THEORY  OF  ELECTROLYTIC  DISSOCIATION? 

Given  the  theory  of  solvation  in  solution  together  with  that  of  elec- 
trolytic dissociation,  the  first  question  that  arises  is,  does  the  former 
really  aid  us  in  explaining  the  phenomena  presented  by  solutions? 

Shortly  after  the  theory  of  electrolytic  dissociation  was  proposed,  it 
was  recognized  and  repeatedly  pointed  out,  that  after  all  it  is  only  a 
theory  of ' '  ideal  solutions,  "i.e.,  very  dilute  solutions.  It  was  shown  not 
to  be  able  to  explain  many  of  the  phenomena  presented  by  even  fairly 
concentrated  solutions.  Indeed,  it  frequently  could  not  deal  quanti- 
tatively with  the  very  solutions  with  which  we  work  in  the  laboratory. 
The  explanation  of  this  shortcoming  was  not  fully  seen,  and  an 
analogy  was  resorted  to.  It  was  pointed  out  that  the  laws  of  Boyle 
and  Gay-Lussac  for  gases  hold  only  for  ''ideal  gases,"  i.  e.,  dilute  gases, 
but  do  not  hold  for  gases  of  any  considerable  concentration. 

It  was  stated  that  the  gas  laws  when  applied  to  solutions  could  not 
be  expected  to  hold  more  generally  than  when  applied  to  gases,  and 
there  the  matter  was  allowed  to  rest. 

l"  Evidence  obtained  in  this  laboratory  during  the  past  twelve  years  for  the  solvate  theory  of 
solution."  Zeit.  phys.  Chem.,  74,  325  (1910). 

2"  Conductivity  and  viscosity  in  mixed  solvents,"  by  H.  C.  Jones  and  co-workers,  Carnegie 
Inst.  Wash.  Pub.  Nos.  80  and  180. 


DISCUSSION   OF  EVIDENCE.  195 

In  the  early  days  of  the  theory  of  electrolytic  dissociation  it  was, 
however,  pointed  out  that  we  have  a  fairly  satisfactory  explanation 
of  why  the  simple  gas  laws  do  not  hold  for  concentrated  gases,  this 
being  expressed  in  the  equation  of  Van  der  Waals;  while  no  analogous 
explanation  was  offered  in  the  case  of  solutions.  That  this  point  was 
well  taken  is  obvious.  A  theory  of  solution,  to  be  of  the  greatest  value, 
must  be  applicable  to  all  solutions,  regardless  of  the  nature  of  the  sub-' 
stance,  regardless  of  the  nature  of  the  solvent,  and  regardless  of  the 
concentration  of  the  solution. 

The  explanation  of  these  apparent  exceptions  to  the  theory  of  elec- 
trolytic dissociation  presented  by  concentrated  solutions  has  been  fur- 
nished by  the  solvate  theory.  We  now  know  that,  for  solutions  in 
general,  a  part  of  the  solvent  is  combined  with  the  dissolved  substance. 
While  the  amount  of  the  solvent  combined  with  any  one  ion  is  greater 
the  more  dilute  the  solution,  at  least  up  to  a  certain  point,  the  total 
amount  of  the  solvent  in  combination  with  the  dissolved  substance  is 
greater  the  more  concentrated  the  solution. 

That  the  amount  of  combined  solvent  may  become  very  great,  even 
relative  to  the  total  amount  of  solvent  present,  can  be  seen  from  the 
following  facts:  In  a  normal  solution  of  calcium  chloride  about  two- 
fifths  of  the  total  water  present  is  combined  with  the  dissolved  sub- 
stance. In  a  three-normal  solution  of  calcium  chloride  about  five- 
sevenths  of  the  total  water  is  combined. 

In  the  case  of  a  normal  solution  of  aluminium  chloride  in  water, 
about  five-eighths  of  the  water  present  is  combined  with  the  dissolved 
substance,  while  in  a  two-normal  solution  about  five-sixths  of  the  water 
present  is  in  a  state  of  combination. 

What  we  suppose  to  be  a  normal  solution  of  calcium  chloride  is,  there- 
fore, more  than  1£  tunes  normal,  while  what  we  suppose  to  be  a  three- 
normal  solution  is  in  reality  more  than  eight  tunes  normal.  In  the 
case  of  aluminium  chloride,  what  we  suppose  to  be  a  normal  solution 
is  more  than  twice  normal,  while  what  we  prepare  as  a  twice  normal 
solution  is  about  twelve  times  normal. 

These  few  facts,  taken  from  thousands  of  a  similar  character,  show 
that  even  fairly  concentrated  solutions  are  much  more  concentrated 
than  we  would  suppose  from  the  method  of  their  preparation;  while 
very  concentrated  solutions  are  many  times  more  concentrated  than, 
without  the  facts  of  solvation,  we  should  be  led  to  expect. 

The  general  conclusion  is  that  even  fairly  concentrated  solutions 
are  much  stronger  than  if  no  solvation  occurred,  and  are  much  more 
concentrated  than  we  are  accustomed  to  consider  from  the  amount  of 
substance  added  to  a  given  volume  of  the  solution — more  or  less  of  the 
water  present  being  in  combination  and  only  the  remainder  playing 
the  role  of  solvent.  Without  the  theory  of  solvation,  we  have  hitherto 
regarded  all  of  the  water  present  as  acting  as  solvent. 


196  DISCUSSION   OF   EVIDENCE. 

We  should,  therefore,  not  expect  the  laws  of  gases  to  apply  to  such 
solutions,  when  we  had  no  idea  what  was  their  concentration.  Now 
that  we  know  their  concentration,  we  find  that  the  laws  of  gases  are 
of  as  general  applicability  to  solutions  as  to  gases,  holding  not  simply 
for  dilute,  but  also  for  concentrated  solutions. 

The  theory  of  electrolytic  dissociation,  supplemented  by  the  theory  of 
solvation,  is,  then,  not  simply  a  theory  of  dilute  or  "ideal"  solutions,  but 
a  theory  of  solutions  in  general. 

DOES  THE  SOLVATE  THEORY  AID  IN  EXPLAINING  THE  FACTS  OF  CHEMISTRY 
IN  GENERAL  AND  OF  PHYSICAL  CHEMISTRY  IN  PARTICULAR? 

To  answer  this  question  at  all  fully  would  lead  us  far  beyond  the 
scope  of  this  monograph.  A  few  facts  bearing  upon  this  question  can, 
however,  be  taken  up.  Take,  for  example,  the  action  of  the  hydrogen 
ion  both  in  the  formation  and  saponification  of  esters.  In  the  presence 
of  the  alcohols  the  hydrogen  ion  accelerates  greatly  the  velocity  with 
which  an  ester  is  formed,  while  in  the  presence  of  water  it  causes  the 
ester  to  break  down  into  the  corresponding  acid  and  alcohol. 

In  terms  of  ordinary  chemical  conceptions  it  is  difficult,  not  to  say 
impossible,  to  interpret  these  reactions,  the  hydrogen  ion  under  one  set 
of  conditions  undoing  what  under  other  conditions  it  effected. 

In  terms  of  the  solvation  theory  these  reactions  admit  of  a  very, 
simple  interpretation.  While  the  hydrogen  ion  is  not  strongly  solvated, 
work  in  this  laboratory  has  shown  that  all  ions  are  more  or  less  solvated. 
In  the  presence  of  alcohol  the  hydrogen  ion  therefore  combines  with  a 
certain  amount  of  this  solvent.  The  hydrogen  ion,  plus  the  alcohol  com- 
bined with  it,  unites  with  the  organic  acid,  forming  complex  alcoholated 
ions  which  then  break  down  yielding  the  ester. 

On  the  other  hand,  the  hydrogen  ion  in  the  presence  of  water  com- 
bines with  a  certain  amount  of  this  solvent.  The  hydrated  hydrogen 
ion,  together  with  the  water  united  with  it,  combines  with  the  ester, 
forming  a  complex  hydrated  ion,  which  then  breaks  down  into  the 
corresponding  acid  and  alcohol  setting  the  hydrogen  free  again.  For 
a  fuller  discussion  of  this  reaction  see  the  paper  by  E.  Emmet  Reid.1 

A  reaction  analogous  to  the  above  is  that  of  hydrogen  ions  on  amides 
in  the  presence  of  water  on  the  one  hand,  and  alcohol  on  the  other  hand. 
In  the  presence  of  water  the  hydrated  hydrogen  ion  combines  with 
the  amide,  forming  a  complex  hydrated  ion  which  then  breaks  down 
yielding  ammonia  and  acid,  the  ammonia,  of  course,  combining  with  the 
acid. 

In  the  presence  of  alcohol  the  alcoholated  hydrogen  ion  combines 
with  the  amide,  forming  a  complex  alcoholated  ion,  which  then  breaks 
down  into  ammonia  and  the  ester  of  the  acid  in  question. 

Hydrogen  ions  in  a  mixture  of  water  and  alcohol,  which  would  con- 
tain both  hydrated  and  alcoholated  hydrogen  ions,  give  both  reactions 

Amer.  Chem.  Journ.,  41,  504  (1909). 


DISCUSSION   OF  EVIDENCE.  197 

simultaneously;  but,  as  Reid  has  pointed  out,  in  the  presence  of  an 
equal  number  of  molecules  of  water  and  alcohol,  the  tendency  of  the 
hydrogen  ion  to  hydrate  is  greater  than  the  tendency  to  form  alcohol- 
ates;  and  under  these  conditions  the  first  reaction  proceeds  much  more 
rapidly  than  the  second.1  A  very  large  number  of  types  of  reactions 
could  be  discussed  illustrating  this  same  point,  i.  e.,  the  value  of  the 
solvate  theory  in  interpreting  chemical  reactions. 

When  we  turn  to  physical  chemical  processes,  the  solvation  of  the 
ions  has  to  be  taken  into  account  at  every  turn.  The  velocities  of  the 
ions  are,  of  course,  a  function  of  the  degree  of  their  solvation;  and  the 
behavior  of  the  ions,  both  chemically  and  physically,  is  a  function  of 
their  velocities.  The  effect  of  dilution,  and  especially  of  temperature 
on  reaction  velocities,  is  largely  a  question  of  the  velocities  of  the  ions 
present,  which,  in  turn,  are  a  function  of  the  degree  of  their  solvation. 

In  determining  the  actual  concentration  of  a  solution,  the  amount 
of  the  solvent  combined  with  the  ions  must  be  taken  into  account,  as 
has  already  been  pointed  out;  and  without  knowing  the  actual  concen- 
trations of  solutions  quantitative  chemistry  would  be  impossible. 

The  solvate  theory  has  thrown  a  flood  of  light  on  the  whole  subject 
of  the  conductivity  of  solutions,  or  the  power  of  the  ions  to  carry  the 
electric  current.  It  has  shown  us  why  the  conductivity  of  lithium  salts 
is  less  than  that  of  sodium  and  potassium,  notwithstanding  the  fact 
that  the  lithium  ion  is  much  smaller  and  lighter  than  the  atom  of  sodium 
or  potassium.  We  now  know  that  the  lithium  ion  is  much  more 
hydrated  than  the  ions  of  these  elements,  and  the  mass  of  the  moving 
ion  is  really  much  greater  in  the  case  of  lithium. 

When  we  come  to  the  temperature  coefficients  of  conductivity,  the 
solvate  theory  has  enabled  us  to  interpret  results  which,  without  its 
aid,  would  be  meaningless.  We  now  know  why  ions  with  the  greater 
hydrating  power  have  the  larger  temperature  coefficients  of  conduc- 
tivity. We  know  why  ions  with  the  same  hydrating  power  have 
approximately  the  same  temperature  coefficients  of  conductivity,  and 
why  dilute  solutions  have  larger  temperature  coefficients  of  conduc- 
tivity than  concentrated  solutions;2  and,  did  space  permit,  we  could 
multiply  examples,  almost  without  limit,  of  the  effect  of  the  solvate 
theory  on  physical  or  general  chemistry. 

WHY  IS  THE  NATURE  OF  SOLUTIONS  OF  SUCH  VITAL  IMPORTANCE  NOT 
ONLY  FOR  CHEMISTRY  BUT  FOR  SCIENCE  IN  GENERAL? 

The  fact  is  well  recognized  that  modern  physical  or  general  chemistry 
has  reached  out  into  nearly  every  branch  of  science,  and  has  had  an 
important  influence  on  many  of  them.  The  question  arises:  Why  is 
this  the  case?  The  answer  is  that  physical  or  general  chemistry  is 
primarily  a  science  of  solutions. 

'Amer.  Chem.  Journ.,  41,  509  (1909).  Ibid.,  35,  445  (1906). 


198  DISCUSSION    OF   EVIDENCE. 

This  answer  may  not  at  first  appear  to  be  self-evident,  but  a  moment's 
thought  will  show  its  general  correctness.  The  whole  science  of  chem- 
istry is  primarily  a  branch  of  the  science  of  solutions  in  the  broad 
sense  of  that  term.  By  solutions  is  meant  not  simply  solutions  in 
liquids  as  the  solvent,  but  solutions  in  gases  and  in  solids  as  well ;  and 
not  simple  solutions  at  ordinary  temperatures,  but  also  at  elevated 
temperatures.  If  we  think  of  chemical  reactions  in  general,  we  will 
realize  what  a  small  percentage  of  them  takes  place  out  of  solution. 
Therefore,  the  nature  of  solutions  is  absolutely  fundamental  for  chem- 
istry. This  applies  not  simply  to  general  chemistry,  including  the 
chemistry  of  carbon,  but  also  to  physiological  chemistry,  which  deals 
almost  entirely  with  solutions  in  one  solvent  or  another. 

When  we  turn  to  physics  we  find  solutions  not  playing  as  prominent 
a  role  as  in  chemistry,  but  nevertheless  coming  in  in  many  places. 
The  primary  cells,  secondary  cells,  electrolysis,  polarization,  diffusion, 
viscosity,  surface-tension,  are  all  phenomena  in  which  the  physicist  is 
interested,  and  all  depend  for  their  existence  upon  solution. 

When  we  turn  to  the  biological  sciences  we  find  that  solution  is 
almost  as  important  as  for  chemistry.  Take  animal  physiology ;  here  we 
have  to  deal  very  largely  with  solution  in  the  broad  sense  of  that  term. 
The  same  remark  applies  to  physiological  botany;  and  solutions  are 
very  important  for  both  animal  and  vegetable  morphology,  especially 
in  their  experimental  developments.  Bacteriology  is  fundamentally 
connected  with  solutions,  and  pharmacology  is  based  upon  solutions 
either  without  or  within  the  body  of  the  animal. 

Solution  in  the  broad  sense  is  as  fundamental  for  geology  as  for 
chemistry.  The  igneous  rocks  were  solutions  of  one  molten  mass  in 
another;  and  sedimentary  deposits  came  down  for  the  most  part  from 
solutions  true  or  colloidal  in  water.  The  minerals  crystallized  out  from 
solutions,  and  solutions  of  various  substances,  such  as  carbon  dioxide, 
are  to-day  weathering  the  rocks  and  continually  changing  the  appear- 
ance of  the  face  of  the  globe. 

An  examination  of  facts  such  as  those  referred  to  above  will  show 
that  the  relation  of  physical  or  general  chemistry  to  solutions  is  the 
prune  reason  why  physical  or  general  chemistry  is  so  closely  related 
to  so  many  other  branches  of  natural  science.  This  alone  would  show 
the  importance  of  a  true  and  comprehensive  theory  of  solutions,  not 
alone  for  physical  or  general  chemistry,  but  for  the  natural  sciences  in 
general. 


BIBLIOGRAPHY. 

It  has  seemed  desirable,  at  the  close  of  this  resumt  of  the  more 
important  lines  of  evidence  bearing  on  the  solvate  theory,  to  give  a 
bibliography  of  the  papers  and  monographs  which  have  been  published 
from  this  laboratory  dealing  directly  and  indirectly  with  the  subject. 

PAPERS. 

1.  JONES  AND  CHAMBERS  :  "  On  some  abnormal  freezing-point  lowerings  produced  by  bromides 

and  chlorides  of  the  alkaline  earths."     Amer.  Chem.  Journ.,  23, 89  (1900). 

2.  CHAMBERS  AND  FRAZEB:  "On  a  minimum  in  the  molecular  lowering  of  the  freezing-point 

of  water,  produced  by  certain  acids  and  salts."     Amer.  Chem.  Journ.,  23,  512  (1900). 

3.  JONES  AND  GETMAN:  "The  lowering  of  the  freezing-point  of  water  produced  by  concen- 

trated solutions  of  certain  electrolytes,  and  the  conductivity  of  such  solutions."      Amer 
Chem.  Journ.,  27,  433  (1902). 

4.  JONES  AND  GETMAN:  "The  molecular  lowering  of  the  freezing-point  of  water  produced  by 

concentrated  solutions  of  certain  electrolytes."     Zeit.  phys.  Chem,  46,  244  (1903). 

5.  JONES  AND  GETMAN:  "A  study  of  the  molecular  lowering  of  the  freezing-point  of  water 

produced  by  concentrated  solutions  of  electrolytes."     Phys.  Rev.,  18,  146  (1904). 

6.  JONES  AND  GETMAN:  "On  the  nature  of  concentrated  solutions  of  electrolytes — hydrates  in 

solution."     Amer.  Chem.  Journ.,  31,  303  (1904). 

7.  JONES  AND  GETMAN:  "  Ueber  das  Vorhandensein  von  Hydraten  in  konzentrierten  wasserigen 

Losungen  von  Elektrolyten."     Zeit.  phys.  Chem.,  49,  385  (1904). 

8.  JONES  AND  GETMAN:  "Ueber  die  Existenz  von  Hydraten  in  konzentrierten  wasserigen 

Losungen  der  Elektrolyte  und  einiger  Nichtelektrolyte."     Ber.  d.  chem.  Gesell.,  37,  1511 
(1904). 

9.  JONES  AND  GETMAN:  "The  existence  of  alcoholates  in  solutions  of  certain  electrolytes  in 

alcohol."     Amer.  Chem.  Journ.,  32,  338  (1904). 

10.  JONES  AND  GETMAN:  "The  existence  of  hydrates  in  solutions  of  certain  non-electrolytes, 

and  the  non-existence  of  hydrates  in  solutions  of  organic  acids."     Amer.  Chem.  Journ., 
32,  308  (1904). 

11.  JONES  AND  BASSETT:  "The  approximate  composition  of  the  hydrates  formed  by  certain 

electrolytes  in  aqueous  solutions  at  different  concentrations."     Amer.  Chem.  Journ., 33, 
534  (1905). 

12.  JONES  AND  BASSETT:  "  Der  Einfluss  der  Temperatur  auf  die  Kristallwassermenge  als Beweis 

fur  die  Theorie  von  den  Hydraten  in  Losung."     Zeit.  phys.  Chem.,  52,  231  (1905). 

13.  JONES  AND  BASSETT:  "The  approximate  composition  of  the  hydrates  formed  by  a  number 

of  electrolytes  in  aqueous  solutions ;  together  with  a  brief  general  discussion  of  the  results 
thus  far  obtained."     Amer.  Chem.  Journ.,  34,  291  (1905). 

14.  JONES:  "L'existence    d'hydrates  dans    les  solutions    aqueuses    d'electrolytes."     Journ. 

Chim.  Phys.,  3,  455  (1905). 

15.  JONES  AND  MCMASTER:  "On  the  formation  of  alcoholates  by  certain  salts  in  solution  in 

methyl  and  ethyl  alcohols."     Amer.  Chem.  Journ.,  35,  316  (1906). 

16.  JONES:  "Die  annahernde  Zusammensetzung  der  Hydrate,  welche  von  verschiedenen  Elek- 

trolyten in  wasseriger  Losung  gebildet  werden."     Zeit.  phys.  Chem.,  55,  385  (1906). 

17.  JONES  AND  UHLER:  "  The  absorption  spectra  of  certain  salts  in  aqueous  solution  as  affected 

by  the  presence  of  certain  other  salts  with  large  hydrating  power."     Amer.  Chem.  Journ., 
37,  126  (1907). 

18.  JONES  AND  UHLER:  "The  absorption  spectra  of  certain  salts  in  non-aqueous  solvents,  as 

affected  by  the  addition  of  water."     Amer.  Chem.  Journ.,  37,  244  (1907). 

19.  JONES  AND  PEARCE:  "Dissociation  as  measured  by  freezing-point  lowering  and  by  con- 

ductivity—bearing  on   the   hydrate   theory.     The   approximate   composition   of  the 
hydrates  formed  by  a  number  of  electrolytes."     Amer.  Chem.  Journ.,  38,  683  (1907). 

20.  JONES  AND  STINE:  "The  effect  of  one  salt  on  the  hydrating  power  of  another  salt  present 

in  the  same  solution."     Amer.  Chem.  Journ.,  39,  313  (1908). 

21.  JONES  AND  ANDERSON:  "The  absorption  spectra  of  neodymium  chloride  and  praseodym- 

ium chloride  in  water,  methyl  alcohol,  ethyl  alcohol,  and  mixtures  of  these  solvents." 
Proceed.  Amer.  Philosoph.  Soc.,  47,  276  (1908). 

22.  JONES  AND  JACOBSON:  "The  conductivity  and  ionization  of  electrolytes  in  aqueous  solu- 

tions as  conditioned  by  temperature,  dilution,  and  hydrolysis."     Amer.  Chem.  Journ.,  40, 

355    /1QQC\ 

23.  JONES:  "The  present  status  of  the  solvate  theory."     Amer.  Chem.  Journ.,  41,  19  (1909). 


200  BIBLIOGRAPHY. 

24.  JONES  AND  ANDERSON:  "The  absorption  spectra  of  solutions  of  a  number  of  salts  in  water, 

in  certain  non-aqueous  solvents,  and  in  mixtures  of  these  solvents  with  water."  Amer. 
Chem.  Journ.,  41,  163  (1909). 

25.  JONES   AND   STRONG :    "Die  Absorptionsspektren  gewisser  Salzlosungen."    Phys.  Zeils., 

10,  499  (1909). 

26.  JONES  AND  STRONG:  "The  absorption  spectra  of  various  salts  in  solution,  and  the  effect  of 

temperature  on  such  spectra."     Amer.  Chem.  Journ.,  43,  37  (1910). 

27.  JONES  AND  STRONG:  "The  absorption  spectra  of  various  potassium,  uranyl,  uranous,  and 

neodymium  salts  in  solution;  and  the  effect  of  temperature  on  the  absorption  spectra  of 
certain  colored  salts  in  solution."  Proceed.  Amer.  Philosoph.  Soc.,  48,  194  (1909). 

28.  JONES  AND  STRONG:  "The  absorption  spectra  of  solutions — a  possible  method  of  detecting 

the  presence  of  intermediate  compounds  in  chemical  reactions."  Amer.  Chem.  Journ., 
43,  224  (1910). 

29.  JONES  AND  STRONG:  "The  absorption  spectra  of  certain  uranyl  and  uranous  compounds." 

Phil.  Mag.,  April  1910. 

30.  JONES  AND  STRONG:  "Spectres  d.  absorption  des  solutions.     Possibilite  d.  une  methode 

pour  determiner  la  presence  de  composis  intermediares  dans  les  reactions  chimiques. 
Journ.  Chim.  Phys.,  8,  131  (1910). 

31.  JONES:  "In  hiesigen  Laboratorium  wahrend  der  vergangenen zwolf  Jahre erhaltene  Anhalts- 

punkte  fur  die  Existenz  von  Solvaten  in  Losung."     Zeit.  phys.  Chem.,  74,  325   (1910). 

32.  JONES  AND  STRONG:  "The  absorption  spectra  of  certain  salts  of  cobalt,  erbium,  neodymium, 

and  uranium,  as  affected  by  temperature  and  by  chemical  reagents."  Amer.  Chem. 
Journ.,  45,  1  (1910). 

33.  JONES:  "Sur  la  position  de  la  theorie  des  solvatea."     Journ.  Chim.  Phys.,  9,  217  (1911). 

34.  JONES  AND  STRONG:  "The  absorption  spectra  of  comparatively  rare  salts.     The  spectro- 

photography  of  certain  chemical  reactions,  and  the  effect  of  high  temperature  on  the 
absorption  spectra  of  non-aqueous  solutions."  Amer.  Chem.  Journ.,  47,  27  (1912). 

35.  JONES:  "The  nature  of  solution."     Journ.  Franklin  Institute,  217  (March  1912). 

36.  JONES:  "Absorption  spectra  and  the  solvate  theory  of  solution."     Phil.  Mag.,  730  (May 

1912). 

37.  JONES:  "  Die  Absorptionsspektra  von  Losungen."     Zeit.  phys.  Chem.,  80, 361  (1912). 

38.  JONES  AND  GUY:  "  Die  Absorptionsspektren  wasseriger  Losungen  von  Neodym-  und  Praseo- 

dymsalzen,  mit  dem  Radiomikrometer  gemessen."     Phys.  Zeit.,  13,  649  (1912). 

39.  JONES  AND  GUY:  "The  absorption  spectra  of  solutions  as  affected  by  temperature  and  by 

dilution.  A  quantitative  study  of  absorption  spectra  by  means  of  the  radiomicrometer." 
Amer.  Chem.  Journ.,  49,  1  (1913). 

40.  GUY,  SHAEFFER,  AND   JONES:   "Die  Anderung  der  Absorption  des  Lichtes  durch  Wasser 

infolge  der  Gegenwart  stark  hydrierter  Salze,  nachgewiesen  mit  Hilfe  des  Radiomikrome- 
ters — ein  neuer  Beweis  fur  die  Solvat-theorie  des  Losungen."  Phys.  Zeit.,  14, 278  (1913) . 
Also.  Amer.  Chem.  Journ.,  49,  265  (1913). 

41.  GUY  AND  JONES:  "The  absorption  spectra  of  a  number  of  salts  as  measured  by  means  of  the 

radiomicrometer."     Amer.  Chem.  Journ.  (November  1913). 

CONDUCTIVITY,  TEMPERATURE  COEFFICIENTS  OF  CONDUCTIVITY,  AND  DISSOCIATION 
IN  AQUEOUS  SOLUTIONS. 

42.  JONES  AND  WEST:  "A  study  of  the  temperature  coefficients  of  conductivity  in  aqueous 

solutions,  and  on  the  effect  of  temperature  on  dissociation."  Amer.  Chem.  Journ.,  34, 
357  (1905). 

43.  JONES  AND  JACOBSON:  See  above,  No.  22. 

44.  JONES  AND  WHITE:  "The  effect  of  temperature  and  dilution  on  the  conductivity  of  organic 

acids  in  aqueous  solution."     Amer.  Chem.  Journ.,  42,  520  (1909). 

45.  CLOVER  AND  JONES:  "The  conductivities,  dissociations,  and  temperature  coefficients  of 

conductivity  between  35°  and  80°  of  solutions  of  a  number  of  salts  and  organic  acids." 
Amer.  Chem.  Journ.,  43,  187  (1910). 

46.  WHITE  AND  JONES:  "The  conductivity  and  dissociation  of  organic  acids  in  aqueous  solu- 

tion at  different  temperatures."     Amer.  Chem.  Journ.,  44,  159  (1910). 

47.  WEST  AND  JONES:  "The  conductivity,  dissociation,  and  temperature  coefficients  of  con- 

ductivity at  35°,  50°,  and  65°  of  aqueous  solutions  of  a  number  of  salts."  Amer.  Chem. 
Journ.,  44,  508  (1910). 

48.  WIGHTMAN  AND  JONES:  "A  study  of  the  conductivity  and  dissociation  of  organic  acids  in 

aqueous  solution  between  zero  and  thirty-five  degrees."  Amer.  Chem.  Journ.,  46,  56 
(1911). 

49.  HOSFORD  AND  JONES:  " The  conductivities,  temperature  coefficients  of  conductivity,  and 

dissociation  of  certain  electrolytes."     Amer.  Chem.  Journ.,  46,  240  (1911). 

50.  WINSTON  AND  JONES:  "The  conductivity,  temperature  coefficients  of  conductivity,  and 

dissociation  of  certain  electrolytes  in  aqueous  solution  from  0°  to  35°.  Probable  induc- 
tive action  in  solution,  and  evidence  for  the  complexity  of  the  ion."  Amer.  Chem.  Journ., 
46,  368  (1911). 


BIBLIOGRAPHY.  201 

51.  WIGHTMAN  AND  JONES:  "A  study  of  the  conductivity  and  dissociation  of  certain  organic 

acids  at  35°,  50°,  and  65°."     Amer.  Chem.  Journ.,  48,  320  (1912). 

52.  SPRINGER  AND  JONES:  "A  study  of  the  conductivity  and  dissociation  of  certain  organic 

acids  in  aqueous  solution  atdifferent  temperatures."    Amer.  Chem.  Journ.,  48, 411  (1912). 

53.  HOWARD  AND  JONES:  "The  conductivity,  temperature  coefficients  of  conductivity,  and 

dissociation  of  certain  electrolytes  in  aqueous  solution  at  35°  50°  and  65°  "  Amer 
Chem.  Journ.,  48,  500  (1912). 

54.  SHAEFFER  AND  JONES:  "A  study  of  the  conductivity,  dissociation,  and  temperature  coeffi- 

cients of  conductivity  of  certain  inorganic  salts  in  aqueous  solution,  as  conditioned  by 
temperature,  dilution,  hydration,  and  hydrolysis."  Amer.  Chem.  Journ.,  49,  207  (1913). 

55.  SMITH  AND  JONES:  "Conductivity,  temperature  coefficients  of  conductivity,  dissociation] 

and  dissociation  constants  of  certain  organic  acids  between  0°  and  65°."  Amer  Chem 
Journ.,  50,  1  (1913). 

56.  JONES:  "  The  bearing  of  hydrates  on  the  temperature  coefficients  of  conductivity  of  aqueous 

solutions."     Amer.  Chem.  Journ.,  35,  445  (1906). 

WORK  IN  MIXED  SOLVENTS. 

57.  JONES  AND  LINDSAY:  "A  study  of  the  conductivity  of  certain  salts  in  water,  methyl,  ethyl, 

and  propyl  alcohols,  and  in  mixtures  of  these  solvents."  Amer.  Chem.  Journ.,  28,  329 
(1902). 

58.  JONES  AND  MURRAY:  "The  association  of  a  liquid  diminished  by  the  presence  of  another 

associated  liquid."     Amer.  Chem.  Journ.,  30,  193  (1903). 

59.  JONES  AND  BASSETT:  "Determination  of  the  relative  velocities  of  the  ions  of  silver  nitrate 

in  mixtures  of  the  alcohols  and  water,  and  on  the  conductivity  of  such  mixtures."  Amer. 
Chem.  Journ.,  32,  409  (1904). 

60.  JONES  AND  CARROLL:  "A  study  of  the  conductivities  of  certain  electrolytes  in  water, 

methyl,  and  ethyl  alcohols,  and  in  mixtures  of  these  solvents.  Relation  between  conduc- 
tivity and  viscosity."  Amer.  Chem.  Joum.,  32,  521  (1904). 

61.  JONES  AND  BINGHAM:  "The  conductivity  and  viscosity  of  solutions  of  certain  salts  in  mix- 

tures of  acetone  with  methyl  alcohol,  with  ethyl  alcohol,  and  water."  Amer.  Chem. 
Journ.,  34,  481  (1905). 

62.  JONES,  LINDSAY,  AND  CARROLL:  "Ueber  die  Leitfahigkeit  gewisser  Salze  in  gemischten 

Losungsmitteln:  Wasser,  Methyl,  Athyl,  und  Prophylakohol."  Zeit.  phys.  Chem.,  56, 
129  (1906). 

63.  JONES  AND  MCMASTER:  "The  conductivity  and  viscosity  of  solutions  of  certain  salts  in 

water,  methyl  alcohol,  ethyl  alcohol,  acetone,  and  binary  mixtures  of  these  solvents." 
Amer.  Chem.  Journ.,  36,  326  (1906). 

64.  JONES  AND  ROUILLER:  "The  relative  migration  velocities  of  the  ions  of  silver  nitrate  in 

water,  methyl  alcohol,  ethyl  alcohol,  and  acetone,  and  in  binary  mixtures  of  these  sol- 
vents, together  with  the  conductivity  of  such  solutions."  Amer.  Chem.  Journ.,  36,  443 
(1906). 

65.  JONES,  BINGHAM,  AND  MCMASTER  :  "Ueber  Leitfahigkeit  und  innere  Reibung  von  Losungen 

gewisser  Salze  in  den  Losungsmittelgemischen,  Wasser,  Methylalkohol,  Athylalkohol,  und 
Aceton."  Zeit.  phys.  Chem.,  57,  193,  257  (1907). 

66.  JONES  AND  VEAZEY:  "A  possible  explanation  of  the  increase  in  viscosity  which  results  when 

the  alcohols  are  mixed  with  water;  and  of  the  negative  viscosity  coefficients  of  certain 
salts  when  dissolved  in  water."  Amer.  Chem.  Journ.,  37,  405  (1907). 

67.  JONES  AND  VEAZEY:  "Die  Leitfahigkeit  und  innere  Reibung  von  Losungen  gewisser  Salze 

in  Wasser,  Methylalkohol,  Athylalkohol,  Aceton,  und  binaren  Gemischen  dieser  Losungs- 
mittel."  Zeit.  phys.  Chem.,  61,  641  (1908). 

68.  JONES  AND  VEAZEY  :  "  Die  Leitfahigkeit  und  innere  Reibung  von  Tetraathylammoniumjpdid 

in  Wasser,  Methylalkohol,  Athylalkohol,  Nitrobenzol,  und  binaren  Gemischen  dieser 
Losungsmittel."  Zeit.  phys.  Chem.,  62,  44  (1908). 

69.  KREIDER  AND  JONES:  "The  dissociation  of  electrolytes  in  non-aqueous  solvents  as  deter- 

mined by  the  conductivity  and  boiling-point  methods."  Amer.  Chem.  Journ.,  45,  282 
(1911). 

70.  JONES  AND  MAHIN:  "The  conductivity  of  solutions  of  lithium  nitrate  in  ternary  mixtures 

of  acetone,  methyl  alcohol,  ethyl  alcohol,  and  water;  together  with  the  viscosity  and 
fluidity  of  these  mixtures."  Amer.  Chem.  Journ.,  41,  433  (1909). 

71.  JONES  AND  MAHIN:  " Conductivity  and  viscosity  of  dilute  solutions  of  lithium  nitrate  and 

cadmium  iodide  in  binary  and  ternary  mixtures  of  acetone  with  methyl  alcohol,  ethyl 
alcohol,  and  water."  Zeit.  phys.  Chem.,  69,  389  (1909). 

72.  SCHMIDT  AND  JONES:  "Conductivity  and  viscosity  in  mixed  solvents  containing  glycerol. 

Amer.  Chem.  Journ.,  42,  37  (1909). 

73.  GUY  AND  JONES:  "Conductivity  and  viscosity  in  mixed  solvents  containing  glycerol. 

Amer.  Chem.  Journ.,  46,  131  (1911). 


202  BIBLIOGRAPHY. 

74.  KREIDEB  AND  JONES:  "The  conductivity  of  certain  salts  in  methyl  and  ethyl  alcohols  at 

high  dilutions."     Amer.  Chem.  Journ.,  46,  574  (1911). 

75.  DAVIS  and  JONES:  " Leitfahigkeits-  und  negative  Viskozitiitskoeffizienten  gewisser  Rubid- 

ium und  Ammoniumsalze  in  Glycerin  und  in  Gemischen  von  Glycerin  mit  Wasser  von 
25°  to  75°."  Zeit.  phys.  Chem.,  81,  68  (1912). 

76.  WIGHTMAN,   DAVIS,  HOLMES,   and  JONES:    "  Conductibilites  et  viscosites   des  solutions 

d'iodure  de  potassium  et  d'iodure  de  sodium  dans  des  melanges  d'alcohol  ethylique 
et  d'eau.  Journ.  Chim.  Phys.,  12,  385  (1914). 

77.  DAVIS,  HUGHES,  AND  JONES:  "Conductivity  and  viscosity  of  solutions  of  rubidium  salts 

in  mixtures  of  acetone  and  water.     Zeit.  phys.  Chem.,  85,  513  (1913). 

78.  JONES:  " Evidence  bearing  on  the  solvate  theory  of  solution."     Journ.  Franklin  Institute. 

479,  677  (Nov.  and  Dec.  1913). 

79.  JONES  and  GUY:  "Eine  quantitative  untersuchung  der  Absorptionsspektren  von  Losungen 

mittels  des  Radiomikrometers."     Ann.  der  Phys.,  43,  555  (1914). 

80.  SHAEFFER,  PAULUS,  and  JONES:  "Die  Anderung  der  Absorption  des  Lichtes  durch  Wasser 

infolge  der  Gegenwart  stark  hydrierter  Salze  gemessen  mit  Hilfe  des  Radiomikrometers," 
Phys.  Zeit.,  15,  447  (1914). 

81.  WIGHTMAN,  WIESEL,  and  JONES:  "A  preliminary  study  of  the  conductivity  of  certain 

organic  acids  in  absolute  ethyl  alcohol."     Journ.  Amer.  Chem.  Soc.,  36,  2243  (1914). 

82.  JONES:  "  Absorptionsspektra  und  die  Solvattheorie  der  Losungen."     Zeit.  Elektrochem.,20, 

552  (1914). 

MONOGRAPHS  ON  RESEARCHES  DEALING  DIRECTLY  OR  INDIRECTLY  WITH 
SOLVATION,  PUBLISHED  BY  THE  CARNEGIE  INSTITUTION  OF  WASHINGTON. 

1.  HYDRATES  IN  AQUEOUS  SOLUTION:  Evidence  for  the  existence  of  hydrates  in  solution,  their 

approximate  composition,  and  certain  spectroscopic  investigations  bearing  upon  the 
hydrate  problem.  By  Harry  C.  Jones,  with  the  assistance  of  F.  H.  Getman,  H.  P. 
Bassett,  L.  McMaster,  and  H.  S.  Uhler.  Carnegie  Institution  of  Washington  Publica- 
tion No.  60  (1907). 

2.  CONDUCTIVITY  AND  VISCOSITY  IN  MIXED  SOLVENTS:  A  study  of  the  conductivity  and  vis- 

cosity of  certain  electrolytes  in  water,  methyl  alcohol,  ethyl  alcohol,  and  acetone,  and 
in  binary  mixtures  of  these  solvents.  By  Harry  C.  Jones  and  C.  F.  Lindsay,  C.  G. 
Carroll,  H.  P.  Bassett,  E.  C.  Bingham,  C.  A.  Rouiller,  L.  McMaster,  and  W.  R.  Veazey. 
Carnegie  Institution  of  Washington  Publication  No.  80  (1907). 

3.  THE  ABSORPTION  SPECTRA  OF  SOLUTIONS  of  certain  salts  of  cobalt,  nickel,  copper,  iron, 

chromium,  neodymium,  praseodymium,  and  erbium  in  water,  methyl  alcohol,  ethyl 
alcohol,  and  acetone,  and  in  mixtures  of  water  with  the  other  solvents.  By  Harry  C. 
Jones  and  John  A.  Anderson.  Carnegie  Institution  of  Washington  Publication  No. 
110  (1909). 

4.  A  STUDY  OF  THE  ABSORPTION  SPECTRA  of  solutions  of  certain  salts  of  potassium,  cobalt, 

nickel,  copper,  chromium,  erbium,  praseodymium,  neodymium,  and  uranium,  as  affected 
by  chemical  agents  and  by  temperature.  By  Harry  C.  Jones  and  W.  W.  Strong. 
Carnegie  Institution  of  Washington  Publication  No.  130  (1910). 

5.  THE  ABSORPTION  SPECTRA  OF  SOLUTIONS  OF  COMPARATIVELY  RARE  SALTS,  including  those 

of  gadolinium,  dysprosium,  and  samarium.  The  spectrophotography  of  certain  chemical 
reactions,  and  the  effect  of  high  temperature  on  the  absorption  spectra  of  non-aqueous 
solutions.  By  Harry  C.  Jones  and  W.  W.  Strong.  Carnegie  Institution  of  Washington 
Publication  No.  160  (1911). 

6.  THE  ELECTRICAL  CONDUCTIVITY,   DISSOCIATION,  AND  TEMPERATURE    COEFFICIENTS  OF 

CONDUCTIVITY  FROM  0°  TO  65°  OF  AQUEOUS  SOLUTIONS  OF  A  NUMBER  OF  SALTS  AND 
ORGANIC  ACIDS.  By  Harry  C.  Jones.  The  experimental  work  by  A.  M.  Clover,  H.  H. 
Hosford,  S.  F.  Howard,  C.  A.  Jacobson,  H.  R.  Kreider,  E.  J.  Shaeffer,  L.  D.  Smith, 
A.  Springer,  Jr.,  A.  P.  West,  G.  F.  White,  E.  P.  Wightman,  and  L.  G.  Winston. 
Carnegie  Institution  of  Washington  Publication  No.  170  (1912). 

7.  THE  FREEZING-POINT,  CONDUCTIVITY,  AND  VISCOSITY  OF  SOLUTIONS  OF  CERTAIN  ELECTRO- 

LYTES IN  WATER,  METHYL  ALCOHOL,  ETHYL  ALCOHOL,  ACETONE,  AND  GLYCEROL,  AND 
IN  MIXTURES  OF  THESE  SOLVENTS  WITH  ONE  ANOTHER.  By  Harry  C.  Jones  and  col- 
laborators. (The  seven  collaborators  in  this  monograph  are  Drs.  C.  M.  Stine,  J.  N. 
Pearce,  H.  R.  Kreider,  E.  G.  Mahin,  M.  R.  Schmidt,  J.  Sam  Guy,  and  P.  B.  Davis.) 
Carnegie  Institution  of  Washington  Publication  No.  180  (1913). 

8.  THE  ABSORPTION  SPECTRA  OF  SOLUTIONS  AS  AFFECTED  BY  TEMPERATURE  AND  BY  DILU- 

TION. A  QUANTITATIVE  STUDY  OF  ABSORPTION  SPECTRA  BY  MEANS  OF  THE  RADIO- 
MICROMETER.  By  Harry  C.  Jones  and  J.  Sam  Guy.  Carnegie  Institution  of  Washington 
Publication  No.  190  (1913). 


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